■ 

I 

I 

I 



^ o = 



a 
a 

o 

W 

CO 




5g 



9 


s 


s 


p 


p 


g 








P. 


CO 


t- 1 


o 

CO 


3 


Pi 




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ffl 



§ 



LESSONS 



IN 



CHEMISTRY. 



BY 

WILLIAM H. GREENE, M.D , 

EMERITUS PROFESSOR OF CHEMISTRY IN THE PHILADELPHIA CENTRAL- 
HIGH SCHOOL, ETC. 



SECOND EDITION 

THOR UGHL Y RE VISED 

BY 

HARRY F. KELLER, Ph.D., 

PROFESSOR OF CHEMISTRY IS THE PHILADELPHIA CENTRAL HIGH SCHOOL. 




PHILADELPHIA '. 

.T. B. LIPPINCOTT COMPANY. 

LONDON : 6 HENRIETTA STREET, COVENT GARDEN. 

1898. 

two cents wcEtVEO 



«=b1 b*s 



QD3 3 



2918 

Copyright, 1884, by J. B. Lippincott Company. 



Copyright, 1898, by J. B. Lippincott Company. 



17° 



TABLE OF CONTENTS. 



LESSON PAGE 

I. — Introduction — Chemical Phenomena 7 

II. — Hydrogen 16 

III. — Oxygen — Combustion 23 

IV. — Composition of Water — Chemical Laws and Theories . . 32 

V. — Laws of Combination — Atomic Theory .... 38 

VI. — Properties of Water — Potable and Mineral Waters . . 45 

VII. — Chemical Nomenclature — Ozone — Hydrogen Dioxide . . 50 

VIII.— Chlorine— Chlorides . ... t ... 57 

IX.— Hydrochloric Acid— Acids— Salts 62 

X.- -Bromine — Iodine — Fluorine 68 

XL — Sulphur — Hydrogen Sulphide 73 

XII. — Sulphur Dioxide — Sulphur Trioxide 79 

XIII.— Sulphuric Acid 82 

XIV.— Sulphates 87 

XV. — Nitrogen — The Atmosphere — Argon 91 

XVI. — Ammonia and its Compounds 97 

XVII. — Ammonium Compounds — Nitrogen Iodide . . . 101 

XVIII.— Nitrous Oxide— Nitric Oxide 104 

XIX. — Nitrogen Peroxide — Nitrogen Pentoxide .... 108 

XX.— Nitric Acid 112 

XXL— Nitrates 116 

XXII. — Phosphorus — Hydrogen Phosphide 119 

XXIII. — Oxides and Acids of Phosphorus 123 

XXIV. — Arsenic — Compounds and Tests 128 

XXV.— Antimony 134 

XXVI.— Boron 137 

XXVII.— Silicon— Glass 140 

XXVIIL— Carbon 144 

XXIX.— Oxides of Carbon 150 

XXX. — Carbonates .........< 156 

XXXI. — Carbon Disulphide — Cyanogen 162 

XXXII. — Hydrocyanic Acid — Cyanides 166 

XXXIII.— Cyanates— Urea . . 171 

XXXIV. — Compounds of Carbon and Hydrogen (1) 

Methane and Saturated Hydrocarbons . . . .175 



4 TABLE OF CONTENTS. 

LESSON PAGE 

XXXV. — Compounds of Carbon and Hydrogen (2) 

Petroleum — Unsaturated Hydrocarbons .... 181 
XXXVI. — Compounds of Carbon and Hydrogen (3) 

Aromatic Hydrocarbons — Elementary Analysis . . 186 

XXXVII.— Methyl and Ethyl Alcohols— Alcoholic Beverages . . 192 

XXXVIII.— Alcohols -Glycols— Glycerol 197 

XXXIX.— Simple Ethers 201 

XL. — Aldehydes — Carbon Acids — Acetone 206 

XLI. — Ethereal Salts — Fatty Acids — Saponification . . .211 
XLII. — Lactic, Oxalic, Tartaric, and Citric Acids . . . .218 

XLIIL— Carbohydrates 220 

XLIV. — Benzene Derivatives, Phenol, Nitrobenzene, Aniline . . 226 
XLV. — Benzene Derivatives, Benzoic, Salicylic, Gallic, and Tan- 
nic Acids — Camphors — Indiyo ..... 230 

XLVI.— Natural Alkaloids 236 

XLVIL — Metals — Spectrum Analysis 241 

XLVIII.— Metallic Compounds— Specific Heat 246 

XLIX. — Lithium — Sodium — Potassium 250 

L.— Silver 257 

LI. — Calcium — Strontium — Barium 265 

LII.— Lead 272 

LIII. — Magnesium — Zinc — Cadmium 277 

LIV.— Copper . . ■ 283 

LV.— Mercury 290 

LVL— Bismuth and Gold 295 

LVIL— Aluminium 301 

LVIII.— Iron and its Metallurgy 307 

LIX. — Compounds of Iron 315 

LX.— Cobalt— Nickel— Manganese 319 

LXL— Chromium and Tin 325 

LXII. — Platinum and its Allied Metals 331 

LXIIL— The Chemistry of Life 334 

Appendix. 

I. — Crystallography 341 

II. — Stereochemistry 346 

Index 349 



PREFACE TO THE SECOND EDITION. 



In preparing the present edition, the aim has been chiefly to 
make such corrections and additions as were rendered necessary 
by the rapid advance of chemical science since the first appear- 
ance of this book. The general plan and arrangement, which 
have proved satisfactory in the experience of the editor as 
well as that of the author, have not been materially modified : 
a few of the chapters have been partly rewritten, and a brief 
explanation of stereoisomerism is given in the Appendix. 

H. F. K. 

Philadelphia, January, 1S98. 



THE DECIMAL SYSTEM OF WEIGHTS AND MEASURES. 
AND THE CENTIGRADE SCALE OF THE THER- 
MOMETER, ARE USED IN THIS BOOK. 



Centigrade Fahrenheit 
Scale. Scale. 



1 Metre = 39.370708 inches. 

I Centimetre = 0.39370 " 
1 Millimetre = 0.03937 



1 Inch 



2.539954 centimetres. 



1 Milligramme = 
1 Centigramme = 
1 Decigramme — 
1 Gramme = 

1 Decagramme = 
1 Hectogramme = 
1 Kilogramme = 



OUNCES TROY 
= 480 GRAINS. 

0.000: >32 
0.000321 
0.003215 
0.032150 
0.321507 
3.215072 
32.150726 



POUNDS 
AVOIRDUPOIS. 
O.0U00022 
0.0000220 
0.0002204 
0022040 
0.0220462 
0.2204621 
2.2046212 



1 Gramme = 15.4323 grains. 



1 Grain = 0.064799 grammes. 

1 Oz. Troy = 31.103496 

1 Lb. Avoirdupois — .453593 kilogrammes. 

1 Cubic Centimetre of water at 4° weighs 1 
gramme. 



To convert centigrade degrees into Fahrenheit 
degrees, multiply by 9, divide by 5, and add 32°. 

To convert Fahrenheit degrees into centigrade 
degrees, subtract 32°, then multiply by 5, and 
divide by 9. 



300° 


572° 


200° 


360° 


150° 


302° 


100° 


212° 


90° 


194° 


80° 


176° 


70° 


158° 


60° 


140° 


50° 


122° 


40° 


104° 


30° 


86° 


20° 


68° 


10° 


50° 


0° 


32° 


—10° 


14° 


—17.8° 


0° 


—20° 


—2° 


—30° 


—22° 


—40° 


—40° 



Water boils. 



Water freezes. 



Mercury freezes. 

A dull red heat is about 500° centi- 
grade, or 950° Fahrenheit. 

A high red heat is about 1000° C, 
and a white heat about 1500°. 



LESSONS IN CHEMISTRY. 



LESSON I. 
INTRODUCTION. 



Chemistry is the science which studies the differences of different 
hinds of matter. 

1. Substance. — Matter occupies space, and can be measured 
and weighed. The different kinds of matter constitute so many 
different substances, which are distinguished from one another by 
general properties, such as color, relative weight, hardness, etc., 
and no two substances can be alike in all properties. 

Some substances are capable of existing in the three possible 
states, as solid, liquid, and gas. Water is the most common 
example of such a substance ; by the action of more or less heat 
it can be converted at pleasure into steam, liquid water, or ice. 
However, if we strongly heat a piece of wood or some sugar, these 
substances will not be melted into liquids or changed to vapor, but 
will be transformed into entirely different kinds of matter, from 
which we cannot again obtain the original substance. 

2. Physical Changes. — We all know that water is capable of 
existing in many forms : mist, fog, rain, frost, snow, sleet, and ice 
all represent the same substance, and we know that these forms 
may change one into another while otherwise the substance 
remains unaffected. The salt water of the ocean differs from the 
fresh water of the rivers which flow into it, only because the sea- 
water contains salt and other forms of matter ; but these substances 
are not water. If salt water be boiled, and a plate or other cold 
surface be held in the steam given off, drops of water will condense 



8 



LESSONS IN CHEMISTRY. 



on the plate, and will be found to be perfectly fresh. The changes 
which convert ice into water, or water into steam or ice, are called 
physical changes, because the nature of the substance is not 
affected ; a little more heat, or a little less, changes the water into 
steam or ice, or the steam or ice back again to water. 

3. Chemical Changes. — Water may, however, undergo other 
changes, in which its nature is altered and new substances are 
produced. Such an alteration in substance is a chemical change. 

Let us fill a test-tube with water, and, closing the open end 
with the thumb, invert it in a small vessel of water. Then we 
wrap a morsel of sodium (see § 414) about as large as a pea, in a 
small piece of wire gauze, twist around this a wire which may 
serve as a handle, and now, raising the tube so that its mouth is 
just below the surface of the water, we push the gauze under the 
edge of the tube (Fig. 1). A small piece of sodium must be used, 




Fig. 1. 



for the experiment is often ended by a little explosion, which 
might break the tube if a large piece were taken. As soon as the 
water touches the sodium, bubbles of gas rise to the surface and 
are collected in the tube. If the latter be not quite filled, the 
gauze may be withdrawn, perfectly dried by holding it in a flame, 
and another piece of sodium introduced, so that the tube shall be 
quite filled with gas. When this is accomplished, we may raise 
the tube from the water, still carefully holding it bottom upwards, 



INTRODUCTION. 



and on introducing into it a lighted match or taper the gas will 
take fire, but will extinguish the taper, which will, however, be 
relighted by the burning gas as it is withdrawn (Fig. 2). 

We shall presently learn that this gas, 
which is called hydrogen, does not come 
from the sodium, nor does it contain any 
sodium ; it must therefore be produced 
from the water, and that portion of the 
latter which has yielded the hydrogen 
must be completely altered in its nature. 
The change is called a chemical reaction. 

4. Elements and Compounds. — If 
water is thus capable of yielding another 
substance produced wholly from the 
water, we must believe that water is a 
compound substance, composed of more 
than one kind of matter. Of the in- 
numerable forms of matter with which 
we are familiar, all, excepting compara- 
tively few, are compounds, and by vari- 
ous means may be converted into simpler 
forms. 

Chemists are acquainted with seventy- 
two substances which they have been unable to change into 
more simple forms. These substances are called elements. 

5. Mercuric oxide is a heavy, red powder. We introduce a 
small quantity of this powder into a test-tube, and heat it in the 
flame of a spirit-lamp or Bunsen burner (Fig. 3). We will first 
notice that its color darkens ; but this change is only physical, for 
if we remove the tube and allow it to cool, the original color is 
restored. If, however, we continue to heat it, in a short time a 
bright mirror forms on the glass in the cooler part of the tube : 
we now light a splint, allow it to burn for a moment, and then 
blow it out, so that it may still retain a spark of fire ; on intro- 
ducing this spark into the tube it at once bursts into flame, and 
the wood is relighted. The experiment may be repeated a num- 




Fig. 2. 



10 



LESSONS IN CHEMISTRY. 



ber of times, extinguishing the splint and relighting it. When 
we have sufficiently studied this phenomenon, we may examine the 
tube, and we will find that the mirror in the interior is composed 
of little globules of metallic mercury ; we may shake them out and 
unite them in one. The gas which has been given off, and which 
causes such brilliant combustion, is called oxygen. The mercuric 

oxide was a compound 
body. By the aid of 
heat we have separated 
it or decomposed it into 
two other substances, 
— mercury and oxygen. 
Mercury and oxygen 
are elements ; chemists 
have not been able to 
convert either of them 
into other substances 
of simpler nature. 

6. Sulphur and cop- 
per are also elements. 
We all know the yellow 
color and brittleness of 
sulphur, and the red 
color and flexibility of 
copper. We will put 
into a test-tube like that used in the last experiment a few small 
pieces of sulphur, and on top of them some copper turnings or a 
bunch of copper wire. We heat the tube ; the sulphur melts, and 
presently begins to boil ; but in a few moments we notice that the 
copper becomes very hot, much hotter than the portion of the 
tube which contains only sulphur. A chemical phenomenon is 
taking place, and the chemical action develops great heat. When 
the experiment has terminated, we allow the tube to cool, break it, 
and find that it no longer contains copper, and unless we have 
used too much sulphur we will find that the latter also has dis- 
appeared. In the place of the sulphur and copper there is a black, 




Fig. 3. 



INTRODUCTION. 11 

brittle substance which resulted from the chemical union of the 
two elements. This substance is called copper sulphide. 

In our first experiment we have seen the decomposition of a 
compound ; in the second we have caused the combination of two 
elements. Combination is the union of two or more substances 
to form a more complex substance, decomposition the separation 
of one substance into more simple substances or into elements. 

If we were to carefully collect and weigh all the oxygen and 
mercury, we would find the weight exactly equal to that of the 
mercuric oxide from which they were obtained. If we make the 
second experiment in a long tube so that no sulphur vapor may 
escape, the weight of the tube will be the same whether it contain 
sulphur and copper or copper sulphide resulting from their union. 
Nothing is lost or gained in either combination or decomposition. 

7. Chemical Combination is not Mixture. — Mercuric oxide 
is not a mixture of oxygen and mercury, nor is copper sulphide a 
mixture of copper and sulphur. We may grind the last two sub- 
stances to the finest powders and mix them together, but in this 
mixture we can by the aid of a microscope distinguish the parti- 
cles of each substance. No microscope would enable us to detect 
sulphur and copper in the black copper sulphide. Since, how- 
ever, we can separate the oxygen from the mercury, as we have 
seen, and the sulphur from the copper, we must believe that the 
elements still exist in their compounds. 

8. Molecules. — We know that under certain conditions any 
substance may change its volume. When a piece of iron is heated 
it grows larger ; when it is cooled it becomes smaller. We can- 
not believe that the matter of the iron actually increases in size 
when it is warmed, although it occupies a greater volume. We 
can understand this change in volume by believing that there are 
in the iron spaces or pores which increase in size when the metal 
expands, and grow smaller when it contracts. These pores must 
be very small, for we cannot perceive them by the aid of the 
most powerful microscopes. That all substances must be porous, 
we can satisfy ourselves by a simple experiment. 

We have a glass tube about a foot long, closed at one end, and 



12 LESSONS IN CHEMISTRY. 

near the other blown out in two bulbs, and the part between the 
bulbs is rather narrow (Fig. 4). We pour water into this tube 
until it is filled to the top of the lower bulb. The water has been 
recently boiled, to drive out the air which was dissolved in it. 
We then fill the remainder of the tube with alcohol, and cork it 
tightly ; the alcohol, which we have colored with a little aniline 
dye, does not at once mix with the water, because the latter is the 
heavier. We now invert the tube, and we see the lighter 
alcohol rise through the water, and at the same time the 
two liquids become thoroughly mixed. But as they mix 
we see little bubbles forming, and there is presently a small 
empty space at the top of the tube. This space is not filled 
with air ; for if we put the mouth of the tube under water 
and draw the cork, the water will rise and fill the tube. 
The mixture then does not occupy as much volume as the 
substances before mixture. We must explain the experi- 
ment by saying that the water and alcohol are porous, and 
that the water runs into the pores of the alcohol, and the 
alcohol into the pores of the water. 

If substances be in this manner porous, they must consist 
of small particles which are separated from one another by 
spaces. Both spaces and particles are so small that we 
can never hope to see them, but we have reason to believe that 
the spaces are quite large in comparison to the particles of matter 
which they separate. These particles are called molecules, and for 
our purposes of study we may consider that the spaces between 
them are perfectly empty. 

9. Atoms and Molecules. — The little particles of which a 
chemical substance is composed are called molecules, and we shall 
learn reasons for believing that all molecules of the same kind, 
that is, of the same substance, have the same size and weight. 

The kind of matter in a molecule must be the same as that in 
any quantity of the substance. If from mercury we can obtain 
only mercury, the molecules of mercury must consist of that ele- 
ment only ; but if from mercuric oxide we can obtain both mer- 
cury and oxygen, the molecules of mercuric oxide must contain 



INTRODUCTION. 13 

both elements. Hence it follows that there must be particles 
still smaller than molecules, and to these smallest particles we 
give the name atoms. Since chemists cannot separate oxygen 
into any other substances, we believe that the atoms of oxygen 
are unalterable by any known force, be it physical or chemical ; 
and it is the same with the atoms of all other elements. 

10. An atom is the ultimate result of the division of matter ; 
it is the smallest particle of an element that can enter into com- 
bination. The nature of an element depends upon the nature 
of its atoms. 

11. A molecule may consist of one or of several atoms ; in the 
latter case, if the atoms be of the same kind the molecule will be 
that of an element or simple body, but if they be of different 
kinds the molecule must be that of a compound. The molecules 
of hydrogen can contain only atoms of hydrogen, and the mole- 
cules of oxygen must consist of atoms of oxygen only, but the 
molecules of water contain atoms of both hydrogen and oxygen ; 
those of copper sulphide contain atoms of sulphur and of copper. 

The nature of a substance will then depend upon the number 
and kind of atoms contained in its molecules. We have seen that 
mercuric oxide contains mercury and oxygen : let us pour a little 
nitric acid on some of this red powder contained in a test-tube, 
and warm the mixture over a lamp. The mercuric oxide disap- 
pears, and we obtain a colorless liquid ; if we pour this liquid into 
a flat dish, and set it aside in a warm place, we will find after a 
time a mass of white crystals. A chemical change has taken 
place : while dissolving in the nitric acid, the mercuric oxide has 
been converted into mercuric nitrate. The latter body contains 
mercury, oxygen, and nitrogen, and its molecule must consist of 
atoms of each of those elements : the new molecule is more com- 
plex than that of mercuric oxide, which contained only two kinds 
of atoms. 

12. Chemical Affinity. — The force which unites atoms is 
called affinity. Its energy is not the same for all atoms, and 
depends on many conditions. While heat may aid in the forma- 
tion of a compound, as in the case of copper sulphide, it may also 
cause decomposition, as in that of mercuric oxide. Other forces, 



14 LESSONS IN CHEMISTRY. 

light and electricity, may act in the same manner, in one case pro- 
ducing combination, and in another decomposition : in every case 
the result depends upon the energy with which the atoms of the 
molecule are held together. Why must we heat the copper and 
sulphur before they will combine ? Simply because the atoms of 
sulphur hold strongly together in the molecules of sulphur, and 
the atoms of copper in the molecules of that metal : we must 
therefore communicate to these molecules so much energy in the 
form of heat that the atoms of sulphur may be sufficiently loosed 
from each other to catch hold of the atoms of copper. Then in- 
stead of molecules containing atoms of sulphur or copper only, we 
have others containing sulphur and copper. 

13. We mix a small quantity of powdered cupric oxide, a black 
compound containing only copper and oxygen, with about one- 
seventh its weight of powdered charcoal, and heat the mixture 
in a test-tube. When the mixture becomes hot, we see that 
the black color changes to reddish brown : after the powder has 
cooled, we turn it out, and find that it is very finely divided 
copper. We cannot heat the cupric oxide alone hot enough to 
decompose it ; but the charcoal, which is an element, has a strong 
affinity for the oxygen, and easily takes it away from the copper. 

The charcoal and oxygen combine together and 
fcj\;\V form a gas called carbon dioxide, which passes 
out of the tube. We can prove that some 
new gas is formed during the experiment, for 
if we push a lighted match into the mouth of 
the tube while it is being heated, the flame 
will be extinguished, an effect exactly oppo- 
site to that which was produced while heat- 
ing the mercuric oxide. We then explain the 
experiment, which is at the same time a com- 
bination and a decomposition, by saying that 
the oxygen has a stronger affinity for the char- 
coal than for the copper. 

14. Into a jar or glass nearly filled with water (Fig. 5) we 
pour first a few drops of a solution of mercuric chloride, and then 




INTRODUCTION. 15 

some solution of potassium iodide. At once a pink or red pre- 
cipitate is formed, showing that a chemical change has taken 
place. Both mercuric chloride and potassium iodide are color- 
less: the first contains two elements, mercury and chlorine, while 
the second also contains two, potassium and iodine. The chemical 
change takes place because the affinities of potassium for chlorine, 
and of mercury for iodine, are stronger than those of potassium 
for iodine, and of mercury for chlorine. Consequently both the 
original substances are decomposed, and two new substances are 
formed, mercuric iodide, which is insoluble (unless too much of 
either substance has been used), and potassium chloride, which 
remains dissolved in the liquid. This is an example of double 
decomposition, the most common kind of chemical change. A 
comparison will show that it closely resembles the double decom- 
position between sulphur molecules and copper molecules already 
explained (§ 12). 

15. Chemical affinity is not to be regarded as a special force, 
but only as one form of energy ; it is manifested between atoms, 
which it holds together in the molecules, just as these molecules 
are held together by the force of cohesion. Affinity depends not 
only on the kind of atoms between which it is exerted, but on the 
temperature : elements which have strong affinities for each other 
at a given temperature may not manifest such affinities at other 
temperatures. 

16. Metals and Non-Metallic Elements. — For convenience 
of study the elements are generally divided into two classes, metals 
and non-metals. The reasons for which an element is considered 
to be a metal or a non-metal will be understood when we shall 
have progressed farther, but we will then learn that the classifi- 
cation is more for convenience than because of any absolutely 
special properties of either class. 

Many of the elements are quite rare, and are seldom seen even 
by chemists ; others are abundant and widely distributed. A 
list of those thus far discovered will be found in the table on 
page 44. 



16 LESSONS IN CHEMISTRY. 

The names of the non-metals are as follows : 



Hydrogen. 


Oxygen. 


Nitrogen. 


Carbon, 


Helium. 


Sulphur. 


Phosphorus. 


Silicon. 


Chlorine. 


Selenium. 


Arsenic. 


Argon. 


Bromine. 


Tellurium. 


Antimony. 




Iodine. 




Boron. 




Fluorine. 









Because they resemble one another in important properties., 
certain of these elements are classed together in natural families. 
We will be better able to understand these relations when we 
have studied some of the compounds, and have seen how all 
chemical changes through which these compounds may pass can 
be explained by our theory of atoms and molecules. At the 
same time we will find that our study will greatly enlarge and 
render more definite the ideas which we have already acquired. 



LESSON II. 

HYDROGEN. 



17. As we have already seen, hydrogen is one of the elements 
of water, of which it constitutes one-ninth by weight. It exists 
in combination with other elements in all animal and vegetable 
substances, in coal, and in the natural oils, petroleum and pitch. 

In our first experiment (§ 3) we have studied one method by 
which it may be obtained, — the action of the metal sodium on 
water. That method is unsuitable for the preparation of any but 
very small quantities of hydrogen ; when it is desired to prepare 
the gas from water, steam may be passed over red-hot iron. 
The metal then combines with the oxygen of the water, setting 
free the hydrogen. 

18. Preparation. — In the laboratory, hydrogen is made by 
the reaction of zinc with hydrochloric acid or sulphuric acid 
diluted with water. 



HYDROGEN. 



17 



We put in the bottom of a tall jar (Fig. 6) some small pieces of 
very thin sheet zinc, or 
a handful of granulated 
zinc, and on this pour 
some hydrochloric acid. 
A brisk effervescence 
begins ; when we apply 
a flame at the mouth of 
the jar, the gas which is 
disengaged at once takes 
fire, and a large stream 
of very pale flame shoots 
into the air. 

This gas is hydrogen. 
Hydrochloric acid is a 
compound of chlorine 
and hydrogen ; when it 
acts on zinc, that metal 
drives the hydrogen out 
of its combination, and 
unites with the chlorine, 
forming a new com- 
pound, called zinc chlo- 
ride. We may express the chemical change as follows : 




Fig. 6. 



BEFORE THE REACTION. 

Hydrochloric acid -f Zinc 

containing 
Hydrogen + Chlorine 



AFTER THE REACTION. 

= Zinc chloride -f Hydrogen 
containing 
Zinc + Chlorine 



Dilute sulphuric acid is usually employed instead of hydro- 
chloric acid for the preparation of hydrogen. We may explain 



the change in a similar manner : 

Sulphuric acid + Zinc = 

containing 
Sulphur + Oxygen + Hydrogen 



Zinc sulphate + Hydrogen 

containing 
Sulphur + Oxygen -f Zinc 



It is sufficient to put the zinc in a bottle, and, after pouring in 
the dilute sulphuric acid, to close the mouth of the bottle with a 

2 



18 



LESSONS IN CHEMISTRY. 




cork through which passes a tube for the exit of the gas ; but it 
is more convenient to have a cork with two holes, or a bottle with 
two necks. Into such a bottle (Fig. 7) we will introduce some 
granulated zinc that has been made by melting 
zinc and pouring it from a little height into a 
bucket of water. Then we adapt to one of the 
necks of the bottle a cork through which passes a 
long tube with a funnel at the upper end ; the 
lower end of this tube must pass nearly to the 
bottom of the bottle, so that it may dip into the 
liquid and no gas may escape by it. To the other 
neck we adapt a cork bearing a tube bent at right 
angles, and this serves for the passage of the gas. 
jjj| Over the end of this tube we may pass a rubber 
pipe and lead the gas wherever we wish it. We 
now pour through the funnel-tube some sulphuric 
acid which we have diluted with about five times 
its volume of water and allowed to cool, for sul- 
phuric acid becomes very hot when it is mixed with water, and 
we always make the mixture by pouring the acid into the water, 
and not the water into the acid. The effervescence shows us that 
gas is being disengaged, and, after waiting a few moments to allow 
the hydrogen time to drive all the air out of the bottle, we may 
make some experiments with our gas. These experiments will 
make us acquainted with its properties. 

19. Properties of Hydrogen. — Hydrogen is a colorless gas, 
and has neither taste nor odor, as we can determine by examining 
it as it escapes from the tube. It is the lightest substance known. 
We connect our gas-generating bottle with the rubber pipe, the 
other end of which is passed over a straight glass tube, and push 
this tube up to the bottom of a wide test-tube which is turned up- 
side down (Fig. 8). In a short time this little jar is filled with 
hydrogen, for the gas is so light that it collects in the jar, and 
pushes the air down and out at the mouth. We can prove that 
the jar is filled with hydrogen, for when we withdraw the tube 
and introduce a lighted taper, the gas at once takes fire and burns 



HYDROGEN. 



19 



r\ 




Fig. 8. 



at the mouth of the jar ; the taper is extinguished on entering the 

gas, but is relighted as it is drawn out through the 

hydrogen flame. The hydrogen is collected in this 

case by upward dry displacement : it displaces the air. 

We again fill our tube with hydrogen in the same 

manner, and taking another and smaller tube we place 

it alongside of the first, which we carefully incline 

(Fig. 9) more and more until we have poured all the 

hydrogen up into the second jar. On introducing 

a lighted taper into the latter, the gas takes fire and 

burns with a slight explosion, for while flowing out of 

the first vessel it became mixed with a little air. 

On account of its lightness, hydrogen is often used 
to fill balloons ; soap-bubbles, which may be easily 
made by dipping the end of the tube into suds, will 
rise quickly in the air when they are shaken from the tube. 

A given volume of hydrogen is only 0.0693 as heavy as the 
same volume of air : this is expressed by saying that the density 
of hydrogen compared to air is 
0.0693 ; for equal volumes, air is 
then 14.44 times as heavy as hydro- 
gen. One litre of hydrogen meas- 
ured at 0° (the freezing point of 
water), and under the ordinary 
pressure of the atmosphere, weighs 
0.0899 of a gramme. 

20. The diffusibility of a gas is 
its tendency to mix with other 
gases : gases will mix with one an- 
other in this manner even through 
the pores of many substances which 
are sensibly porous, that is, possess 
pores large enough to be seen by the aid of a microscope. It 
has been found that the diffusibility of gases depends upon their 
densities. The lighter a gas is, the more diffusible is it also, and, 
on the contrary, the heavier gases do not diffuse as quickly as the 




Fig. 9. 



20 



LESSONS IN CHEMISTRY. 




Fig. 10. 



lighter ones. Since hydrogen is the lightest gas, we can under- 
stand that it must be the most diffusible : we allow a little hydro- 
gen to escape into the air, and in a few seconds it scatters through 
all the air in the room. 

We have arranged another tube through which the hydrogen 
may escape from our gas-bottle, and this tube is drawn out so that 
it has a small opening at which we may 
burn the gas if we desire. Close above 
the unlighted gas escaping from this jet 
we hold a piece of paper (Fig. 10) ; the 
hydrogen passes through the paper, as we 
prove by igniting it above, and the flame 
of the gas quickly sets fire to the paper 
and passes through to the gas at the jet. 

Because it is so diffusible, hydrogen 
cannot be kept in bottles which have the 
smallest cracks. It even passes through hot plates of iron and 
platinum. 

21. Hydrogen is almost insoluble in water, and may be col- 
lected over the pneumatic trough by wet displacement. Gases 
which are but slightly soluble in water may be collected in 

this manner: the jar in which we 
wish to receive the £as is filled 
with water and inverted in a 
trough near the top of which is 
a shelf on which the jar may rest. 
The water will not run out of the 
jar as long as the mouth of the 
latter is below the surface. Under 
the edge of this jar filled with 
water we pass the end of the tube 
from which escapes the gas that 
we wish to collect ; this gas bub- 
bles up through the water, which it drives out of the jar (Fig. 
11). If it be desired, we can transfer the gas from one jar to 
another, by first filling the second jar with water, placing it on 
the trough, and then pouring the gas up through the water by 




Fig. 11. 




HYDROGEN. 21 

inclining the jar which contains it under the edge of that which 
is to receive it. 

22. Hydrogen is the only gas which conducts heat. We have 
fitted to the ends of a glass tube (Fig. 12) two tightly- fitting corks 
through each of which 
passes a smaller tube, 
and also a thick wire 
which is connected with 
a voltaic battery ; the jp IG# 12. 

two wires are joined by 

a thin platinum wire which becomes heated red-hot by the elec- 
tric current. We can now pass any gas through this tube and 
notice the effect on the wire : we try oxygen, nitrogen, carbon 
dioxide, and see that the wire still remains red-hot ; but when 
we pass hydrogen through the tube, the wire ceases to glow. 
The hydrogen has conducted away the heat. On account of 
its conducting power, and because it has the property of being 
absorbed by certain metals, we believe that hydrogen is a sort 
of metallic vapor. By very great pressure and extreme cold, 
hydrogen has been converted into a liquid. 

23. We have seen (§ 19) that hydrogen will burn in air, and 
that it will not support combustion ; the burning taper was ex- 
tinguished by hydrogen. When it burns, hydrogen combines 
with the oxygen of the air, forming vapor of water ; this is the 
sole product of the combustion of hydrogen. We may assure 
ourselves of this by holding over a jet of burning hydrogen an 
inverted jar, of which the interior will rapidly become covered 
with little drops of dew, and these will soon unite together and 
trickle down the sides of the jar. This takes place with hydro- 
gen which has been perfectly dried by passing through a tube 
containing calcium chloride, or pumice-stone wet with sulphuric 
acid (Fig. 13) ; both these substances remove all moisture from 
gases with which they come in contact. 

If hydrogen be mixed with half its volume of oxygen, or about 
three times its volume of air, the mixture will explode violently 
when ignited. For this reason we must be careful that all the 
air has been driven from the generating bottle before lighting the 



LESSONS IN CHEMISTRY. 



hydrogen. We may make the explosion harmlessly by passing a 
little hydrogen into a hydrogen pistol, made of sheet tin, and, after 





Fig. 13. 

corking the mouth of the pistol, ignite the mixed gases by holding 
a flame to the little hole at the other end ; 
the cork is then driven out with a loud 
noise (Fig. 14) : while charging the pistol, 
we must close the hole with the finger. 
If we slip over a small jet of burning 

hydrogen a rather wide glass tube (Fig. 15), we will find that 
when the flame has reached a certain point 
in the wide tube it begins to quiver, and a 
more or less musical tone is produced. The 
tone may be varied by using tubes of differ- 
ent lengths : it is caused by the current of 
air ascending the tube. 

24. Certain very finely divided metals 
have the power of absorbing hydrogen so 
rapidly as to become hot enough to light the 
gas. Spongy platinum is such a substance ; 
when a small piece of this very porous form 
of platinum, tied by a thin wire in the centre 
of a small brass ring (Fig. 16), is held in a 

jet of escaping hydrogen, it becomes bright hot and the gas is 

inflamed. The spongy platinum should be heated shortly before 




Fig 



OXYGEN. — COMBUSTION. 23 

making the experiment. It is not hard to understand this phe- 
nomenon, for when the platinum absorbs 
the hydrogen the gas is necessarily re- 
duced to a small volume in the pores of 
the metal, and the heat which keeps the 
molecules of the gas at large distances 
from each other must raise the tempera- 
ture when those distances are diminished 
by the condensation ; just as the heat 
which converts water into steam reap- Fig. 16. 

pears when the steam is condensed. 

Hydrogen combines with many of the other elements, but 
under ordinaiy circumstances its affinities are not very pro- 
nounced. Heat is required to bring about the union of hydro- 
gen and oxygen. Hydrogen and chlorine combine under the 
influence of light (see § 73). Pure hydrogen is not poisonous, 
but it does not support respiration (see § 33). 




LESSON III. 
OXYGEN.— COMBUSTION. 

25, Oxygen was discovered by Scheele and independently by 
Priestley in 1774. It is the most abundant element at the surface 
of the earth ; it forms about one-fifth of the atmosphere, in which 
it exists uncombined, but mixed with the element nitrogen and 
other gases ; its combination with hydrogen is water, and it enters 
largely into the composition of nearly all minerals and rocks. 

We have seen (§ 5) that oxygen is produced when mercuric 
oxide is heated ; but this method would be too expensive for the 
preparation of large quantities of oxygen. 

26. Preparation. — The most convenient process for obtaining 
oxygen consists in heating a compound of chlorine, oxygen, and 
potassium, called potassium chlorate. This is a white, crystalline 



24 



LESSONS IN CHEMISTRY. 



substance, from which heat drives out all the oxygen, leaving a 
compound of potassium and chlorine, called potassium chloride. 

We put a little potassium chlorate in a test-tube, and heat it 
rather strongly in the flame of a spirit-lamp or Bunsen burner. 

It melts, and soon begins to 
boil ; this boiling is the es- 
cape of the oxygen, as we 
can prove by pushing into 
the tube a match-stick bear- 
ing a spark of fire, which 
instantly bursts into flame 
(Fig. 17). The white sub- 
stance which remains in the 
tube after all the oxygen 
is driven out, is potassium 
chloride. 

When we wish to make 
and collect larger quantities 
of oxygen, we mix the po- 
tassium chlorate with about 
one-eighth its weight of 
manganese dioxide, which 
causes the gas to be given 
off at a lower temperature, and with less danger of explosion. The 
manganese dioxide is a black powder, and is found unaltered after 
the experiment, being simply mixed with the potassium chloride : 
it helps the reaction because it has an affinity for oxygen ; but 
this affinity is so feeble that the new compound which is formed 
is at once decomposed, the oxygen being given off, while the man- 
ganese dioxide remains as it was at first. We may consider that 
it pulls the oxygen away from the potassium chlorate. 

We introduce our mixture of potassium chlorate and manganese 
dioxide into a glass flask, to which we adapt a tightly-fitting cork 
bearing a tube for the exit of the gas. Then we place the flask 
on some dry sand in a little tin or sheet-iron dish, which we call 
a sand-bath, and under this we place a lamp. The sand becomes 




Fig. 17. 



OXYGEN.— COMBUSTION. 



25 



hot and heats the flask gradually, and generally prevents cracking 
of the glass. We may now slip a rubber tube over the delivery- 




Fig. 18. 

tube of the flask, and when oxygen begins to come off, as we may 
ascertain by holding a lighted match near the end of the tube, we 
may collect the gas in a jar over the pneumatic trough (Fig. 18). 

As glass flasks often 
break in this experiment, 
when we want many litres 
of oxygen, we heat the 
generating mixture in a 
flask made of tightly 
lapped sheet copper or 
tin plate. As little parti- 
cles of manganese dioxide 
are carried out with the 
gas, we usually wash the 
latter by making it pass 
through some water in a wash-bottle 
shown in Fig. 19. 

When all the oxygen has been disengaged, we remove the end 
of the tube from the water in the trough before taking the heat 
from under the flask : otherwise water would be drawn back as the 
retort cools, and would break a glass flask, and the steam might 
burst one of tin or copper. This precaution is observed in the 
preparation of all gases made by the aid of heat. 




Fig. 19. 
The whole apparatus is 



26 



LESSONS IN CHEMISTRY. 



After filling several jars with oxygen, we remove them from the 
trough by passing a saucer under the mouth of each, below the 
surface of the water ; then on lifting them out, the water in the 
saucer prevents the escape of the gas. We can now turn them 
quickly mouth upward, still keeping covered with the saucer, and 
we are ready to study the gas.* 

27. Properties. — Oxygen has neither color, taste, nor odor. 
When freshly made from potassium chlorate it usually has a 
smoky appearance and more or less odor, but these are impurities, 
and disappear after the gas has stood for a time over the pneumatic 
trough. It is a little heavier than the air, its density being 
1.1056 ; one litre of the gas at 0°, and normal pressure, weighs 
1.437 grammes. It is almost insoluble in water. It has been 
converted into a liquid by great cold and pressure. 

Oxygen manifests energetic affinity for most of the other ele- 
ments : it combines with some of them at ordinary temperatures, 
and with others by the aid of more or less heat. 

28. Combustion. — The burning of wood, coal, illuminating 
gas, oil, and other substances with which we are familiar, is only 
the combination of those bodies with the oxygen of the air. Into 
a small tube closed at one end, we have put some ferrous oxalate, 
made by adding oxalic acid to a solution of ferrous sulphate, and 
after drawing out the open end of the tube so as to leave a small 
thin opening, we twisted a wire about the tube, and heated it 
until no more gas escaped at the opening ; we then sealed the thin 

end of the tube by holding it 
for a moment in the flame. 
The result of this heating has 
been to decompose the ferrous 
Will, oxalate, leaving a very fine 
powder of iron in the tube. 
We now break this tube, and 
shake out the powder, which instantly takes fire, falling in a shower 
of sparks (Fig. 20), if our tube has been well prepared. The iron 

* Oxygen gas compressed in steel bottles can now be procured from dealers 
in chemicals: it is conveniently drawn from these bottles as required for ex- 
periment. 




Fig. 20. 



OXYGEN. — COMBUSTION. 27 

lias combined with the oxygen of the air: it has been burned into 
a substance called iron oxide. Usually it is necessary to heat a 
combustible substance before it will burn ; then as soon as the 
union with oxygen, or the oxidation as we call it, begins, the 
chemical action develops sufficient heat to keep the temperature so 
high that the combustion may continue without further aid. 

Only one-fifth of the air is oxygen, and we shall learn that the 
other gases with which that oxygen is mixed not only do not help, 
but prevent combustion : pure oxygen should then support com- 
bustion much more energetically than air, and we have seen that 
oxygen causes a spark on a match-stick or a taper to burst into 
flame. 

29. We wrap a copper wire around a piece of charcoal, and 
fasten the other end of the wire in a hole in a piece of tin plate 
large enough to cover the mouth of one of our jars. After hold- 
ing the charcoal in a flame until a corner of it becomes red-hot, 
we quickly remove the saucer from the jar and plunge into it the 
charcoal, which remains suspended (Fig. 21). Instantly the com- 
bustion grows very vivid, and, if we have a knotty piece of char- 
coal, brilliant sparks are thrown off. The charcoal 
combines with the oxygen until all of one or the other 
is used up. The result of the combination is a gas 
called carbon dioxide, and if we put a lighted taper 
into the jar containing it, the flame will be extin- 
guished : we may add that the oxygen also has been 
burned, and can serve for no other combustion as long 
as it remains combined with the carbon. 

After softening a steel watch-spring by heating it in 
a flame, we twist it into a coil, one end of which we fasten in a 
cork, and over the other end we slip the split end of a piece of 
match-stick; after lighting this we quickly introduce it into a 
bottle of oxygen (Fig. 22). The flame heats the iron so hot that 
it can begin to burn, and the oxidation furnishes heat enough for 
the combustion to continue : brilliant stars of burning steel shoot 
out, and hot drops of iron oxide fall to the bottom of the jar, in 
which it is well to leave a layer of water to prevent breaking. 




28 



LESSONS IN CHEMISTRY. 




We have prepared a deflagratiDg-spoon by fastening a small 
saucer-like piece of sheet copper on the end of a straight copper 
wire. We support this in the hole in our 
tin plate, and on a little dry sand which we 
put in the spoon we place a piece of phos- 
phorus (§ 177) a little larger than a pea. 
We light the phosphorus — it takes fire very 
easily — and plunge the spoon into a new 
jar of oxygen (Fig. 23). At once a most 
intense light is produced by the combustion 
gggil of the phosphorus, and the jar becomes 
filled with a white smoke of phosphoric 
Jig. 22. oxide, the compound of phosphorus and 

oxygen. After a time this smoke dissolves in the layer of water, 
which we leave in the jar for this experiment as for the last. 

The metal magnesium burns very bril- 
liantly in the air : we twist together half 
a dozen ribbons of this metal, and, after 
fastening one end in our jar-cover, we 
light the other with a match. On intro- 
ducing this into a jar of oxygen the in- 
tensity of the combustion is dazzling. A 
white smoke of magnesium oxide soon 
settles in the jar, and contains of course 
the magnesium and oxygen which have 
combined together. The jar is often 
broken in this experiment. 

These experiments have been only intense cases of what we 
commonly call combustion, a phenomenon which we apply for the 
production of heat and light. The combustible substances ordi- 
narily employed, such as wood, coal, illuminating gas, wax, tallow, 
oil, etc., contain carbon and hydrogen ; charcoal is almost wholly 
carbon ; these substances burn because the carbon and hydrogen 
which they contain, unite readily with the oxygen of the air when 
the union is started by the aid of heat. 

30. The brightness of the light is not always proportional to 




Fig. 23. 



OXYGEN. — COMBUSTION. 29 

the amount of heat : we have seen that the flame of hydrogen is 
very pale, but it is very hot. If we desire to increase the heat of 
a fire, we furnish the combustible with more oxygen by blowing 
air into it with a bellows, and we rake the ashes from our coals in 
order that the oxygen may come in contact with the hot carbon : 
it is possible, however, to furnish too much air, if the latter be 
cold, as we see when we extinguish a candle-flame by blowing on 
it. When we want the most intense combustion possible, we 
supply the burning body with pure oxygen, and the hottest flame 
which we can obtain is that of hydrogen burning in oxygen. This 
flame is that of what is called the oxyhydrogen blow-pipe, in which 
a tube through which the oxygen is forced passes inside of another 
tube carrying the hydrogen ; the two gases, coming from separate 
gas-holders, or caoutchouc bags, mix at the opening of the jet (Fig. 
24). If they were mixed before the moment of burning, the ap- 



Oxygen. 




Hydrogen 



paratus containing the mixture would be burst by the explosive 
union of the gases (§ 23). In using the oxyhydrogen blow-pipe, 
we first turn on the hydrogen, light it, and then slowly turn on the 
oxygen until we have the hottest flame. If it be inconvenient to 
use hydrogen, we may substitute for it illuminating gas, connect- 
ing by a rubber tube the oxyhydrogen blow-pipe with a gas-fix- 
ture. While the oxyhydrogen flame is not very bright, it is very 
hot, and when we hold in it a piece of watch-spring, or an old 
penknife-blade, the iron is burned, making a brilliant fountain of 
fire. The metal platinum, which does not melt at the highest 
furnace heat, melts readily in the oxyhydrogen flame. 

31. Fire is the combustion with incandescence — that is, pro- 
duction of light and heat at the same time — of a solid substance : 



30 LESSONS IN CHEMISTRY. 

we have seen such phenomena in the combustion of charcoal and 
iron. The oxidation takes place only on the surface of the burn- 
ing body. 

32. Flame is the combustion with incandescence of a gas or 
vapor, as in the burning of hydrogen, phosphorus, and magne- 
sium : at the borders of the flame the gas or vapor may mix with 
the air ; but the interior of the flame must consist of highly- 
heated, yet unburned gas. Why are certain flames very bright, 
while others give little or no light ? We burn a little sulphur in 
a deflagrating-spoon in a jar of oxygen, and the combustion, 
though very brilliant, would not serve for illumination. In order 
to produce a bright white light, a flame must contain particles of 
solid matter which may become highly heated. The burning 
phosphorus and magnesium were brilliantly luminous because the 
little solid particles of phosphoric oxide and magnesium oxide 
which were formed, became very hot. The flames of hydrogen 
and sulphur do not give white light because the products of com- 
bustion, water in one case and sulphurous oxide in the other, are 
gases at the high temperature at which they are formed, and 
gases cannot be heated hot enough to give white light. How- 
ever, the products of combustion of tallow, wax, and illuminating 
gas are not solid, yet these substances are useful for artificial light. 
In these cases the illumination is due to little particles of carbon. 
The combustible gases and vapors come in contact with enough 
oxygen to completely burn them only on the outer edge of the 
flame ; but the heat is radiated into the flame as well as from it, 
and the gases, which are compounds of hydrogen and carbon, are 
decomposed ; little solid particles of carbon are set free, and these 
become very hot and give out light : when they reach the outside 
of the flame they are entirely consumed, unless there be too little 
oxygen, and in that case the flame smokes. When a cold body — 
a piece of glass will answer — is held for a moment in the brightest 
part of a lamp- or gas flame, the little particles of carbon are de- 
posited on the cold surface in the form of soot. 

We may by a very simple means render the colorless flame of 
hydrogen quite brilliant : we have fitted to a bottle a cork through 



OXYGEN. — COMBUSTION. 



31 




which pass two tubes, the outer ends being drawn out to fine jets. 
One of these tubes is short, and passes only through the cork ; 
the other passes to the bottom of the bottle, in which we have 
placed some broken pumice-stone saturated with benzene (Fig. 25). 
To this same tube is joined a short side-tube, which 
we connect with a bottle containing zinc and dilute 
sulphuric acid. When all the air is expelled from 
the bottle, we light the hydrogen at the two jets. 
At one it burns with a colorless flame : it is the 
flame of hydrogen just as it comes from the gen- 
erating bottle. The other flame is quite bright ; the 
hydrogen which has passed through the benzene has 
become charged with the vapor of that volatile liquid, 
and as that vapor, containing hydrogen and carbon, * IG * 2o ' 
is decomposed by the heat before it burns, the carbon particles 
become incandescent. 

When the flame of illuminating gas or of a lamp is supplied 
with oxygen at the inside, the particles of carbon are burned in- 
stantly and do not become hot : the flame then gives no light. 
This is the case in the Bunsen burner (Fig. 26), in which the 
force of the escaping gas draws air through 
holes in a tube surrounding the jet; the air 
and gas mix together, and all the carbon is 
consumed before it can become incandescent. 
We then have a flame which gives great heat, 
but does not deposit smoke on any vessels 
which we may heat in it. 

If we hold a piece of lime in the flame of 
the oxyhydrogen blow-pipe, it becomes very 
hot and emits a brilliant light. This consti- 
tutes the calcium or oxyhydrogen light which 
is used in theatres. Lime is used because it 
is neither burned, melted, nor changed into 
vapor by the intense heat. 

33. Slow Combustion. — All the examples of oxidation which 
we have so far considered are said to be cases of rapid com- 




Fig. 26. 



32 LESSONS IN CHEMISTRY. 

bustion : they take place with the production of intense heat and 
light. Sometimes, however, there is no bright light, no high 
temperature, and yet combustion takes place as certainly as before. 
A piece of iron which rusts by exposure to damp air is only com- 
bining with oxygen, and the rust is a compound of iron with the 
oxygen and moisture of the atmosphere : here the heat of chemi- 
cal union is developed so slowly that it is conducted away by the 
air and surrounding bodies, and the iron does not become heated. 
Respiration is a slow combustion. The warmth of our bodies, 
and all our animal motions, are due to the gradual oxidation of the 
carbon and hydrogen of our tissues. At every breath fresh oxy- 
gen is introduced into the lungs, where it is absorbed by the blood 
and carried through the arteries to the most remote parts of 
the system ; then, when all the oxygen in the blood is used, the 
water and carbon dioxide produced by the combustion are carried 
through the veins to the lungs, and thrown out with the exhaled 
air. Animal life itself depends on this slow oxidation : we all 
know how quickly any animal perishes from suffocation when the 
supply of air is entirely cut off. The muscles of our bodies con- 
tain no force except that which is produced by the combustion of 
their own substance. Great muscular exertion consequently re- 
quires increased oxidation, and we quickly become fatigued when 
we are obliged to burn up our tissues more rapidly than they are 
remade from our food. Also, the quantity and kind of food re- 
quired depend upon the amount and kind of work which we 
must perform. 



LESSON IV. 



COMPOSITION OF WATER.— CHEMICAL LAWS AND 

THEORIES. 

34. Water is the sole product of the combustion of hydrogen 
in air or oxygen. Its composition, that is, the proportion in 
which the hydrogen and oxygen are combined together, may be 



COMPOSITION OF WATER. 



33 



determined by analysis and by synthesis. Analysis is the separa- 
tion and weighing of the constituents of a compound ; synthesis 
means the formation of a substance by causing its elements to 
unite in the proper proportion. 

35. Electrolysis of Water. — Electrolysis means the decom- 
position of a substance by an electric current. For the decomposi- 
tion of pure water an enormously strong current would be re- 
quired, and because we do not desire to use such a strong current 
we employ dilute sulphuric acid : the final result is the same as if 
we were to use water, the sulphuric acid being found unchanged 
after the experiment. Instead of using the acid, we might make 
a strong solution of ordinary salt ; the salt would make the water 
a better conductor of electricity. We introduce the dilute sul- 
phuric acid (about five parts of water to one of acid) into a vessel 
through a hole in the bottom of which are cemented two wires, 
the inner ends of each being soldered to a little plate of thin plat- 
inum. We fill two small test-tubes with water, and, closing the 
mouths with the fingers, invert one over each of these plates : we 
now connect the ends of the wires with the poles of a voltaic bat- 
tery (Fig. 27). Little bubbles of gas at once begin to rise in the 
tubes, and as soon as 
the quantities of gas col- 
lected are large enough 
to allow us to notice the 
volume of each, we see 
that in one of the tubes 
there is twice as much 
as in the other. When 
that tube is filled, we 
raise it carefully, and 
the introduction of a 
lighted match will con- 
vince us that the gas is 
hydrogen. When we raise the other tube, keeping the end closed 
with the thumb, until we are ready to push into it a match-stick 
bearing a spark, the kindling of the spark into flame shows us 

3 




Fig. 27 



34 



LESSONS IN CHEMISTRY. 



that the second gas is oxygen. Water is then composed of two 

volumes of hydrogen combined with one volume of oxygen. 
36. We have seen that the density of hydrogen compared to 

air is 0.0693, and that the density of oxygen is 1.1056. A given 

volume of oxygen must then be 

1.1056 -5- 0.0693 = 16 (a very little less), 

sixteen times as heavy as an equal volume of hydrogen. As we 
have two volumes of hydrogen and only one of oxygen, 
the oxygen in water must weigh eight times as much as 
the hydrogen. 

37. Eudiometric Synthesis of Water. — A eudio- 
meter is a graduated strong glass tube, closed at one end 
near which two thin platinum wires are soldered into 
the glass on opposite sides ; an electric spark may be 
passed between the wires on the inside of the tube 
(Fig. 28). If we fill such a tube with mercury, and, 
after inverting it in a vessel of mercury, pass into it 
some hydrogen, and then half as much oxygen, an 
electric spark will cause the gases to combine ; after the 
little explosion which takes place in the tube, the water 
which is formed is condensed in minute drops in the 
cold tube, and the atmospheric pressure forces the mer- 
cury up, filling the tube completely. Here, again, we 
see that water is composed of two volumes of hydrogen 
combined with one volume of oxygen. 




Fig. 28 Fig. 29. 

38. Synthesis by Weight. — We may make the synthesis of 



CHEMICAL LAWS AND THEORIES. 35 

water by a very instructive method which was first adopted by the 
French chemist Dumas. We prepare hydrogen from sulphuric acid 
and zinc in the ordinary manner, and thoroughly dry it by passage 
through a tube (A) containing little pieces of pumice-stone wet 
with strong sulphuric acid (Fig. 29). We then cause it to pass 
through a tube containing some cupric oxide (B), a black com- 
pound of copper and oxygen, and this tube is connected with a 
U-shaped tube (C) filled with pumice-stone also moistened with 
strong sulphuric acid. The U tube is placed in a vessel contain- 
ing some broken ice. Before connecting our tubes together, we 
have carefully weighed that holding the cupric oxide, and the last 
U tube with its contents. When this whole apparatus is filled 
with the hydrogen coming from the bottle, we heat the cupric 
oxide by a spirit-lamp, and when it becomes hot the hydrogen gas 
takes away the oxygen from the copper. Steam is formed and is 
condensed in the U tube (C). When the color of the cupric oxide 
has entirely changed to red, we warm the whole length of the 
tube containing it, in order to drive all of the water over into the 
U tube : we allow our apparatus to cool, take it apart, and again 
weigh the tubes of which we had determined the weight before 
the experiment. The copper which is left in the first will weigh 
just as much less than the cupric oxide as the latter has lost oxy- 
gen. The increased weight of the U tube (C) will be the weight 
of the water formed, and by subtracting from this weight the 
weight of the oxygen, we will have the weight of the hydrogen 
contained in that water. We find that there is almost exactly 
eight times as much oxygen as hydrogen. In very accurate ex- 
periments we would perfectly purify our hydrogen and adopt all 
possible precautions that no vapor of water might escape from the 
tube C. 

CHEMICAL LAWS AND THEORIES. 

39. No matter by what process water may be formed, no matter 
by what process its composition may be determined, it is always 
found to contain the same proportions of oxygen and hydrogen ; 
never more nor less than eight (7.94 exactly) parts by weight of 



36 



LESSONS IN CHEMISTRY. 



the first for one of the second. If we try to combine the gases 
in other proportions, the excess of the one or other, out of the 
proportion one to eight, will be left uncombined. The analysis 
of all known substances has shown a similar constancy of com- 
position, a constancy which is expressed in the following 

40. Law of Definite Proportions: The proportion by 
weight in which the elements exist in any compound, is invaria- 
ble. This is generally called Dalton's first law. 

41. We have already found that the proportions by volume 
according to which oxygen and hydrogen unite are one to two. 
This is a simple relation of volumes. Experiments with other 
gases will in time show us that when gases combine, there is 
always some such simple relation between the volumes of the gases 
that enter into combination. Thus, 

One volume of hydrogen combines with exactly one volume of chlorine. 
Two volumes of hydrogen combine with exactly one volume of oxygen. 
Three volumes of hydrogen combine with exactly one volume of nitrogen. 
Two volumes of nitrogen combine with exactly one volume of oxygen. 

We might find many more such examples, and the statement 
of these facts constitutes 

Gay-Ltjssac's First Law : there is a simple relation between 
the volumes of gases which combine. 

42. Let us study the volume of the compound 
formed when that compound is in the same con- 
dition as the original elements ; that is, the gaseous 
state. 

By grinding together with emery, we have 
accurately fitted together the necks of two glass 
bottles that have exactly the same capacity (Fig. 
30). We fill the lower one with perfectly dry 
chlorine gas (§ 71), and the upper with dry hydro- 
gen, and then hermetically join them together by 
the ground joint. All of this must be done in a 
room lighted only by a candle or small gas-flame. 
We now allow the apparatus to stand for a day in 
a room where the sunlight may not shine on it 




Fig. 30. 



I 



COMPOSITION OF WATER. 



37 



directly. The gases will slowly combine, and the yellowish color of 

the chlorine will disappear. When we open the bottles under the 

surface of mercury, we will find that no gas escapes from them, and 

no mercury enters. The gas hydrochloric acid has been formed, and 

_ , -,-i_,. -j m. l • f one volume of hydrogen and 

Two volumes of hydrochloric acid must contain < " ° 

I one volume of chlorine. 

Over the open end of a eudiometer we have passed a piece of 
strong rubber tubing, to the other end of which is attached a 
straight glass tube about the size and length of the eudiometer ; 
the rubber tube is firmly tied at each joint. We fill our apparatus 
with mercury, place it vertically in the mer- 
cury trough, and introduce five cubic centi- 
metres of oxygen and ten cubic centimetres 
of hydrogen. We now close the end of the 
tube with the finger, lift it from the trough, 
and, after bringing the two tubes parallel to 
each other, clamp them in that position in a 
stand (Fig. 31). We have tightly fitted a 
perforated cork around the lower end of the 
eudiometer, and on this cork we now as tightly 
fit the lower end of a wide glass tube, which 
we slip over the eudiometer, whose wires we 
have connected with long copper wires that 
pass out at the upper end of the wide tube. 
We now fill the wide tube with perfectly clear 
oil (lard oil or sweet oil) heated to 130° ; we 
adjust the mercury at the same level in the 
two tubes, and, after carefully reading the vol- 
ume occupied by the mixed gases, we pass an electric spark in the 
eudiometer. The gases of course combine : steam is formed, but 
does not condense, because the eudiometer is heated. On again 
making the mercury levels the same, and examining the volume 
of this steam, we find that it is only two-thirds as great as that 
of the mixed gases : in other words, 




Fig. 31. 



Two volumes of steam contain 



two volumes of hydrogen and 
one volume of oxygen. 



38 LESSONS IN CHEMISTRY. 

Ammonia gas is a compound of hydrogen and nitrogen, and its 
analysis proves that two volumes of the gas may be decomposed 
into one volume of nitrogen and three volumes of hydrogen. 

On comparing these results, we find that 

Two volumes of hydrochloric acid contain one volume of hydrogen and one 
volume of chlorine. 

Two volumes of vapor of water contain two volumes of hydrogen and one vol- 
ume of oxygen. 

Two volumes of ammonia contain three volumes of hydrogen and one volume 
of nitrogen. 

We see that not only is there a simple relation between the 
volumes of gases which combine, but, as is expressed in 

Gay-Lussac's Second Law, there is a simple relation between 
the volume of a compound gas and the sum of the volumes of the 
gases which form that compound. 



LESSON V. 

CHEMICAL LAWS AND THEORIES (Continued). 

43. Equivalent Combining Proportions. — Careful analysis 
of hydrochloric acid has shown that it is composed of 35.5 parts 
by weight of chlorine combined with one part by weight of hydro- 
gen. Chlorine combines with mercury, forming a compound 
called mercuric chloride or corrosive sublimate; this compound 
contains for every 35.5 parts of chlorine, exactly 100 parts of 
mercury. We dissolve 135.5 grammes of mercuric chloride in 
water, and put some zinc into the solution : the chlorine has a 
stronger affinity for the zinc than for the mercury ; it conse- 
quently combines with the zinc, forming zinc chloride, which re- 
mains in the solution, while mercury separates. If we wait until 
all the 35.5 grammes of chlorine which were combined with 100 
grammes of mercury have united with the zinc, and then deter- 
mine the quantity of zinc which is required to combine with that 
quantity of chlorine, we would find that the zinc chloride formed 



CHEMICAL LAWS AND THEORIES. 39 

weighs 68.25 grammes : that isj 32.75 (68.25 — 35.5) grammes 
of zinc combine with 35.5 grammes of chlorine. Consequently, as 
far as combining with chlorine is concerned, 32.75 parts of zinc 
have just as much power as 100 parts of mercury. 

Oxygen combines with mercury and with hydrogen : it com- 
bines also with zinc and with chlorine. Analysis of the com- 
pounds so formed shows that 35.5 parts of chlorine will combine 
with 8 parts of oxygen, and that 8 parts of oxygen will combine 
with 32.75 parts of zinc or with 100 of mercury. We have 
already seen (§ 36) that 8 parts of oxygen combine with one part 
of hydrogen. These numbers must then express the relations 
between the combining quantities of the corresponding elements. 
Thousands of analyses have shown that similar equivalent propor- 
tions exist for all of the elements. The combining proportions so 
found must bear simple relations to the relative weights of the 
Qtoms ; for if atoms have a real existence, chemical combination 
must result from the union of one, two, or more atoms of one 
element with one, two, or more atoms of another ; and, since 
combination is in definite proportions, the same substance must 
always result from the union of the same kind of atoms in the 
same proportion. 

44. If it be possible to determine the relative weights of the 
molecules, compared with any unit, and to arrive at definite con- 
clusions as to the number of atoms these molecules contain, then 
we can determine the relative weights of the atoms. Some new 
considerations will enable us to make such determinations. 

LAW OF AVOGADRO AND AMPERE.— ATOMIC THEORY. 

45 Different solid and liquid substances expand in very dif- 
ferent degrees by the action of the same temperature. Gases, 
however, all expand alike. If at the same pressure we raise or 
lower the temperature of equal volumes of different gases through 
the same number of degrees, we find that they all expand or con- 
tract precisely the same proportion of the volume. If expansion 
be separation of molecules from one another (§8), it follows that 
equal volumes of gases, measured at the same temperature and 



40 LESSONS IN CHEMISTRY. 

pressure, contain the same number of molecules. This hypothesis, 
proposed by Avogadro and Ampere, is true if there be such things 
as molecules, and if there be no molecules we can explain no 
chemical phenomena. 

46. If equal volumes of gases contain the same number of 
molecules, the relative weights of those equal volumes must also 
express the relative weights of the molecules. The relative weights 
are the densities, and these densities are usually calculated to ex- 
press the relation between the weight of the gas and that of an 
equal volume of air, which is taken as unity. Since hydrogen is 
the lightest gas, its molecule must have the least weight: the 
density of oxygen is sixteen times as great as that of hydrogen, and 
the molecule of oxygen must be sixteen times as heavy as that of 
hydrogen. Because of the lightness of the molecule of hydrogen, 
chemists have chosen that molecule as the standard of comparison 
for other molecules, — i.e., the molecular weights are referred to it. 
The density of hydrogen compared to air being 0.0693, the air 
is 14.44 times as heavy as hydrogen : consequently if we know 
the density of a gas compared to air, we may easily calculate its 
density compared to hydrogen by multiplying the first by 14.44. 
Thus, the density of vapor of water compared to air is 0.622, com- 
pared to hydrogen it is 0.622 X 14.44 = 9. The molecule of 
water (steam) must then be nine times as heavy as that of hydrogen. 

47. Let us see now whether we can determine the relations 
between the weight of any atom and that of a molecule of hydro- 
gen. According to the law of Avogadro, the oxygen molecule is 
sixteen times heavier than the hydrogen molecule. One volume 
of oxygen combines with two volumes of hydrogen, and the result 
is two volumes of water vapor. How many atoms of hydrogen 
are contained in one molecule of this vapor? To answer this 
question we must remember that two volumes of hydrochloric 
acid, containing just as many molecules as two volumes of steam, 
contain only half as much hydrogen as the two volumes of steam 
(§ 41). If, then, the molecule of hydrochloric acid contains one 
atom of hydrogen (and it cannot contain less) the molecule of 
water vapor must contain two. And if the molecule of water 



CHEMICAL LAWS AND THEORIES. 41 

vapor contains two atoms of hydrogen, the weight of this mole- 
cule must be eighteen times as great as that of an atom of hydro- 
gen, for water contains oxygen and hydrogen in the ratio 8 : 1 
= 16:2. Then, if a molecule of water contains but one atom 
of oxygen, and there is no reason to believe that it contains 
more than one, the oxygen atom must weigh sixteen times as 
much as that of hydrogen. 

48. Now we may apply our theory to the facts already studied. 

Two volumes of hydrogen represent two atoms, each of which 
weighs one : two volumes of oxygen represent two atoms, each of 
which weighs sixteen. Water is formed by the union of two 
atoms of hydrogen and one atom of oxygen, and a molecule of 
water weighs eighteen times as much as an atom of hydrogen. 

A molecule of hydrochloric acid contains one atom of hydrogen 
and one atom of chlorine, and this molecule weighs 36.5 if one 
atom of hydrogen weighs 1. 

A molecule of ammonia contains one atom of nitrogen (weigh- 
ing 14) and three atoms of hydrogen, and is 17 times as heavy as 
an atom of hydrogen. 

But the density of water compared to hydrogen is 9 ; that of 
hydrochloric acid, 18.25, and that of ammonia, 8.5. We see then 
that if an atom of hydrogen occupies one volume and iceighs one, 
the molecule of any substance in a state of gas or vapor must 
occupy twice as much volume as one atom of hydrogen, and the 
weight of the molecule will be expressed by twice the density of the 
gas or vapor referred to hydrogen. In other words, the standard 
of molecular weight is 2, since there are two atoms in the mole- 
cule of hydrogen. 

Different methods are employed for determining the atomic 
weights of the elements. At this point we need only understand 
that if the element be gaseous or volatile, and if we have reason to 
believe that its molecule contains two atoms, then, since the mole- 
cule of hydrogen consists of two atoms, and equal volumes of 
gases contain equal numbers of molecules, the same figures which 
express the density of the gas compared to hydrogen, will express 
also the atomic weight. 



42 LESSONS IN CHEMISTRY. 

This atomic theoiy, which has been slowly developed during the 
present century, furnishes an intelligible explanation of chemical 
phenomena. New discoveries are continually bringing new facts to 
its support, and, though it may be modified by the results of future 
researches, its principal features will probably remain undisturbed. 

CHEMICAL NOTATION. 

49. In order that the composition of a substance, that is, the 
number and kinds of atoms in its molecules, may be understood at 
a glance, we employ a special method of representing elements and 
compounds. The first letter, or the first and another letter, of 
the name of an element, is used to express one atom of that ele- 
ment. Thus, H means one atom of hydrogen ; O, one of oxygen ; 
S, one of sulphur ; C, one of carbon, and Ca, one of calcium. 
These are called the symbols of trie elements. More than one 
atom is expressed by a little figure placed to the right of the 
symbol, slightly above or below its central line ; H 2 or H 2 (read 
H two) means two atoms of hydrogen ; O 4 represents four atoms 
of oxygen. Compounds are then written so that the symbols 
entering into the formula express the number and kind of atoms 
in a molecule. H 2 means a molecule of water, composed of two 
atoms of hydrogen and one of oxygen : H 2 S0 4 means a molecule 
of sulphuric acid, containing two atoms of hydrogen, one of sul- 
phur, and four of oxygen. To express any number of molecules 
we use an ordinary figure placed to the left of the formula ; thus, 
2HC1 means two molecules of hydrochloric acid, each of which 
contains one atom of hydrogen and one of chlorine. 

50. We may now study the molecular changes which have oc- 
curred in the chemical phenomena that we have already observed. 
In the decomposition of water by sodium, one atom of hydrogen 
in each molecule of water is replaced by sodium, and when a mole- 
cule of hydrogen is set free, two molecules of a compound called 
sodium hydroxide are formed. We represent the change thus : 

2H20 + Na 2 = 2NaOH + H 2 

2 molecules of 2 atoms of 2 molecules of 1 molecule of 

water. sodium. sodium hydroxide. hydrogen. 

This chemical equation expresses the changes which take place 



CHEiMICAL LAWS AND THEORIES. 43 

in the chemical reaction. As the symbol for each atom means a 
definite quantity of matter, and as there can be no change in the 
quantity of matter during the reaction, there must be as many 
atoms represented in one member of the equation as in the other. 
When we know what is formed by the reaction of certain mole- 
cules, our equation will enable us to calculate the quantities of the 
substances. The weight of one atom of sodium being 23 ; one 
atom of hydrogen, 1 ; and one atom of oxygen, 16, we find that 
46 grammes of sodium will yield 2 grammes of hydrogen, and 80 
grammes of sodium hydroxide. 

2HOH + Xa 2 = 2XaOH + H 2 

2 (1 + 16 + 1) 23 + 23 2(23 +16 + 1) 1 -f 1. 

We can calculate the volume of the hydrogen at 0°, from its 
weight (§ 19), and we can so estimate the weight and volume 
of hydrogen which will be set free by any given quantity of 
sodium. Chemical formulae and equations thus represent more 
than mere theory ; they exactly express the combining propor- 
tions. 

The action of sulphuric acid on zinc, which yields zinc sulphate 
and hydrogen, is written 

Zn + H 2 SO* = ZnSO* + H 2 
Sulphuric acid. Zinc sulphate. 

Of course we must know by experiment what is formed in a 
reaction, before we can write the chemical equation ; we must also 
know by analysis the proportions of the elements in any compound 
before we can write a formula which we believe to express the 
atoms in its molecule. 

The decomposition of potassium chlorate by heat (§ 26) is 
2KC10 3 = 2KC1 + 30 2 

Potassium chlorate, Potassium chloride, Oxygen, 
2 molecules. 2 molecules. 3 molecules. 

That of mercuric oxide, in the same manner, is 
2HgO = 2Hg + O 2 
Mercuric oxide. Mercury. 

The reaction of hydrogen with cupric oxide, which enabled us 
to make the synthesis of water, is written 

CuO + H 2 = H 2 + Cu 
Cupric oxide. Copper. 



44 



LESSONS IN CHEMISTRY. 



51. The following table gives the names and symbols of the 
elements which are at present known, and the weights of the 
atoms compared to the weight of an atom of hydrogen. Some of 
these atomic weights might be more exactly expressed ; an atom 
of oxygen is 15.88 times as heavy as that of hydrogen ; the exact 
atomic weight of nitrogen is 13.93 ; but these numbers are so 
nearly 16 and 14, that for memory's sake it is preferable to use 
the nearest whole numbers. 



Names of the Ele- 


on 


o4 


Names of the Ele- 


DO 


o 42 

S "tic 


ments. 


a 


*- "a> 


ments. 


a 


•£> "S 




a? 


<zz 




GO 


<* 


Aluminium . . . 


Al 


27 


Mercury (hydrar- 






Antimony (stibium) 


Sb 


120 


gyrum) . . . 


Hg 


200 


Argon 


A 


40(?) 


Molybdenum . . 


Mo 


96 


Arsenic 




As 


75 


Mckel 


Ni 


59 


Barium . . 




Ba 


137 


Niobium .... 


Nb 


94 


Bismuth . . 




Bi 


208 


Nitrogen .... 


N 


14 


Boron . . . 




B 


11 


Osmium .... 


Os 


190 


Bromine . . 




Br 


80 


Oxygen .... 





16 


Cadmium . . 




Cd 


112 


Palladium . . . 


Pd 


106.6 


Caesium . . 




Cs 


133 


Phosphorus . . . 


P 


31 


Calcium . . 




Ca 


40 


Platinum .... 


Ft 


193.5 


Carbon . . 




C 


12 


Potassium (kalium) 


K 


39.1 


Cerium . . 




Ce 


141.2 


Rhodium .... 


Rh 


102.2 


Chlorine . . 




CI 


35.5 


Rubidium . . 


Rb 


85.2 


Chromium 




Cr 


52.5 


Ruthenium . . . 


Ru 


100.9 


Cobalt . . . 




Co 


59 


Scandium . . . 


Sc 


44 


Copper . . 




Cu 


63.1 


Selenium .... 


Se 


79.5 


Didymium . 




Di 


145.4 


Silicon .... 


Si 


28 


Erbium . . 




Er 


166 


Silver (argentum) . 


Ag 


108 


Fluorine . . 




F 


19 


Sodium (natrium) . 


Na 


23 


Gallium . . 




Ga 


69 


Strontium . . . 


Sr 


87.5 


Germanium . 




Ge 


72 


Sulphur .... 


S 


32 


Glucinuin 




Gl 


9 


Tantalum . . . 


Ta 


182 


Gold (aurum) 




Au 


195.7 


Tellurium . . . 


Te 


125 


Helium . . 




He 


4(?) 


Thallium .... 


Tl 


204 


Holmium . . 




Ho 


162 (?) 


Thorium .... 


Th 


234 


Hydrogen 




11 


1 


Tin (stannum) . . 


Sn 


118 


Indium . 




In 


113.4 


Titanium . . . 


Ti 


48 


Iodine . . . 




I 


127 


Thulium .... 


Tu 


170.4(?) 


Iridium . . . 




Ir 


192 


Tungsten (wolfra- 






Iron (ferrum) . 




Fe 


56 


mium) .... 


W 


184 


Lanthanum . . 




La 


139 


Uranium .... 


Ur 


238 


Lead (plumbum 




Pb 


207 


Vanadium . . . 


V 


51 


Lithium . . . 




Li 


7 


Ytterbium . . 


Yt 


173 


Magnesium . . 




Mg 


21 


Yttrium .... 


Y 


89 


Manganese . . 




Mn 


55 


Zinc 


Zn 


65 








Zirconium . . . 


Zr 


90 

- 



PROPERTIES OF WATER. 



45 



LESSON VI. 



PROPERTIES OF WATER.— POTABLE AND MINERAL 

WATERS. 

52. Pure water is not met with in nature. When we desire it, 
it must be prepared by distilling water ; that is, boiling it in a re- 
tort, and condensing the steam. We usually conduct this distilla- 
tion in tin-lined copper retorts. For most of the distillations in 
the laboratory we employ a flask or retort, connected with a long 
tube which is surrounded by a wider tube, and a stream of cold 




Fig. 32. 

water continually flows through the space between the two tubes 
in this condenser, as we call it (Fig. 32). 

Pure water has neither taste nor odor ; although it is colorless 
in small quantity, it has a deep-blue color when in large masses. 
It solidifies when sufficiently cooled, and, since it is always con- 
verted into ice at the same temperature, that temperature is taken 
as 0° in the centigrade thermometer scale which we use in the 
laboratory. 

53. The temperature of water does not change while it is freez- 
ing, and that of ice does not change while it is melting. This is 
because all the heat which is communicated to ice during its 
melting is required to produce the change of state ; indeed, one 



46 LESSONS IN CHEMISTRY. 

kilogramme of ice at 0° requires as much heat to melt it as would 
raise 79 kilogrammes of water from 0° to 1°, or one kilogramme 
from 0° to 79°, and yet the water from the melted ice still has a 
temperature of 0°. Ice is crystallized ; it consists of a great many 
little six-sided pyramids dovetailed together. We can notice the 




ifiG. 



crystalline form of water by examining some snow-flakes that have 
fallen on black cloth (Fig. 33). 

During the cooling of water, it contracts in volume until its 
temperature reaches 4° ; it then begins to expand, and on freezing 
expands considerably. Ice is only 0.93 as heavy as water at 4°. 
Strong vessels are broken by the freezing of water in them, and 
it is the same expansion which kills delicate plants by frost, for 
the ice formed in them tears apart the fibres and destroys the sap- 
vessels. Since it is easy everywhere to obtain water at its point 
of maximum density, that is, 4°, this density has been chosen as 
the unit of density or specific gravity for liquids and solids. It 
is also at this temperature that one litre of water weighs one kilo- 
gramme. 

Water and ice continually emit invisible vapor, but water does 
not begin to boil until its tension of vapor * is equal to the at- 
mospheric pressure. We consider that the normal atmospheric 
pressure is equal to that of a column of mercury 760 millimetres 



* The tension of vapor of a liquid at any temperature is measured by the 
decrease in the height of the mercury in a barometer-tube, up into which the 
liquid is passed in small quantities until no more of it changes into vapor. 
The number of millimetres through which the level of the mercury has then 
fallen, expresses the tension of the vapor. 



PROPERTIES OF WATER. 47 

in height, and under this pressure the boiling point of water is 
selected as the 100° point of the centigrade thermometer. 

While water is boiling, its temperature does not rise : after it 
has reached the boiling point, all the heat passes into the steam, 
where it is required to hold apart the molecules. To convert one 
kilogramme of water at 100° into steam requires as much heat 
as would raise the temperature of 537 kilogrammes of water 
from 0° to 1°, or 5.37 kilogrammes from 0° to 100°. The 
conversion of water into steam expands it 1696 times: that is 3 
one litre of water will yield 1696 litres of steam at 100°. 

54. Chemical Properties. — We have seen that water is de- 
composed by an electric current : it is also decomposed by very 
high temperatures (1200°). We will find that it enters into 
many chemical reactions, in some of which it is decomposed and 
part of its hydrogen set free, as in the experiment with sodium 
(§ 3). In other cases both the oxygen and hydrogen atoms are 
taken into new combinations. Water forms a large proportion of 
animal and vegetable tissues. 

Water dissolves many substances, solid, liquid, and gaseous. 
We all know that salt, sugar, and alum will dissolve in water, be- 
coming for the time part of the liquid. We immerse the bulb of 
a thermometer in a vessel of water, into which we throw some 
ammonium nitrate, and stir the liquid : at once the thermometer 
indicates a lower temperature. When solids dissolve in water, 
cold is produced, because the heat required to separate the mole- 
cules of the solid must be taken from the water. On the con- 
trary, when gases dissolve in water, the liquid becomes warmer, 
because the heat no longer required to hold apart the molecules 
of the gas can now raise the temperature. 

55. Water exerts a very curious action on some substances. 
We take some large blue crystals of cupric sulphate and heat 
them on a piece of tin over a lamp. We see that they gradually 
become white, and crumble into a powder. We throw some of 
this powder into water, and the water becomes blue ; cupric sul- 
phate can only exist in crystals when it is combined with water, 
and it is blue only when combined with water. In the same 



48 LESSONS IN CHEMISTRY. 

manner water is necessary to the crystalline form, and often to 
the color, of many substances, and when combined in this way is 
called water of crystallization. Water of crystallization is chem- 
ically combined in the crystals, for it is always in definite propor- 
tions. In a piece of crystallized cupric sulphate there are five 
molecules of water for every molecule of copper sulphate. 

56. Natural Water. — As it occurs naturally, water always 
contains foreign matters suspended or dissolved in it. These sub- 
stances are derived from the air through which the rain falls, or 
from the soil over which the water flows. 

57. According to the kind and quantity of these matters 
present, the water is potable, mineral, or unfit for drinking 
and cooking. Water used for domestic purposes, such as 
drinking, cooking, and washing, should be cool, limpid, and 
odorless, having a very feeble but pleasant taste that should 
be neither bitter, salty, nor sweet, and should make suds with 
soap without forming a curd. Water which possesses these prop- 
erties always holds in solution a certain quantity of the gases of 
the air, oxygen, nitrogen, and carbon dioxide, and usually a small 
quantity of mineral matters. The gases are absolutely essential 
to good water, but their quantity varies considerably in different 
waters, and at different times in water from the same source. 
This dissolved air separates from water which stands in a warm 
place, and part of it collects on the sides of the vessel in small 
bubbles, which we have all seen in a glass of water that has stood 
for several hours. Fish cannot live in water containing no dis- 
solved oxygen ; they do not breathe, but their gills remove the 
dissolved oxygen from the water which they continually draw 
through those organs (see § 33). 

The solid matters in a potable water should not exceed two or 
three decigrammes per litre, and these matters should be entirely 
mineral. They usually consist of compounds of calcium and mag- 
nesium ; magnesium sulphate and calcium sulphate being the 
most common, while a small proportion of common salt is gener- 
ally present. 

58. When larger quantities of calcium or magnesium com- 
pounds are present, the water is said to be hard. The hardness 



MINERAL WATERS. 49 

may be due either to sulphates or carbonates. Water dissolves 
only a very small quantity of calcium sulphate, but then has a 
peculiar taste and curdles the soap when we use it for washing. 
We add a few drops of a solution of barium chloride to some 
water containing a little calcium sulphate, and instantly a white 
cloud appears : this is caused by the formation of an insoluble 
body called barium sulphate, and the test makes us sure that the 
water contained a sulphate. Calcium and magnesium carbonate 
are insoluble in pure water, but dissolve in water containing car- 
bonic acid gas, or carbon dioxide. When such water is boiled, the 
carbon dioxide is driven out, and then the carbonate separates, for 
it is no longer soluble. Hence we have a method of curing hard 
water which contains only calcium carbonate and carbonic acid : 
we boil it, and allow it to settle, and after pouring off the clear 
water expose it to the air for a time, so that it may dissolve some 
of the gases from the atmosphere. 

59. Drinking-water must not contain animal or vegetable sub- 
stances : they render it very unwholesome. Happily, the waters 
of rivers, which become contaminated with so many such impuri- 
ties, generally become purified during their exposure to the air, 
because the foul matter is gradually oxidized. Water containing 
these matters usually has a sweetish taste and a disagreeable odor, 
which may, however, be very faint. It may be purified by passing 
it through a charcoal filter (§ 227). 

60. Mineral waters contain various dissolved mineral mat- 
ters. Some are hot, others warm, and still others cold. Those 
which effervesce or sparkle contain a considerable proportion of 
carbonic acid gas in solution, and it is the escape of this gas which 
produces the sparkling. Apollinaris water contains, besides the 
carbonic acid gas, principally a little sodium acid carbonate, com- 
mon salt, and magnesium and sodium sulphates. Buffalo lithia 
water contains very little of the sodium compounds, but consider- 
able quantities of calcium sulphate, with carbonates of potassium, 
calcium, barium, and lithium. 

Saratoga water has a large proportion of calcium and magnesium 
carbonates dissolved by the excess of carbon dioxide which it con- 

4 



50 LESSONS IN CHEMISTRY. 

tains, and a very large proportion of common salt. Gettysburg 
water contains principally the carbonates of calcium, magnesium, 
and sodium, together with a little dissolved silica : its excess of 
carbon dioxide is quite small. Hunyadi Janos contains sulphates 
of magnesium, sodium, calcium, and potassium ; these substances 
give to it purgative properties, in which it is resembled by Fried- 
richshall water, for the composition of the latter is somewhat 
similar. 

Chalybeate waters are such as contain either iron carbonate, 
held in solution by an excess of carbon dioxide, or ferrous sul- 
phate : in the former case the water becomes muddy on exposure 
to the air, for as the carbon dioxide escapes, ferrous carbonate is 
deposited. The Mercer County water, of Virginia, contains a 
large proportion of ferrous sulphate. Iron waters are usually 
cold. 

Sulphur waters owe their odors and their virtues to hydrogen 
sulphide and sulphides of potassium and sodium. They are gen- 
erally warm, or even hot. 



LESSON VII. 



NOMENCLATURE OF COMPOUNDS OF OXYGEN- 
OZONE— HYDROGEN DIOXIDE. 

61. Besides being able to express the composition of molecules 
by chemical formulae, as we have learned, it is important that we 
may have a distinctive name for each substance, and that this 
name may express as far as possible the composition of the mole- 
cule. A compound of only two elements is called a binary com- 
pound ; one containing three is a ternary compound ; one con- 
taining four, a quaternary. We may be satisfied at present to 
study a system of naming — a nomenclature — for the binary com- 
pounds of oxygen. These are called oxides. 



NOMENCLATURE OF COMPOUNDS OF OXYGEN, ETC. 



51 




Fig. 34. 



We place a small piece of phosphorus on a piece of glass on a 
plate, light it, and cover it with a bell-jar (Fig. 34). The phos- 
phorus combines with the oxy- 
gen of the air, and the com- 
pound which is formed settles 
like flakes of snow in the jar 
and on the plate. This is an 
oxide of phosphorus. On another 
plate, in the same manner, we 
burn a small piece of sodium : 
we have here formed sodium 
oxide. Now we rinse out each 
jar and plate with a little water : 
when the water comes in con- 
tact with the oxides that have 
been formed, there is a hissing 
noise, and the jars become 

warm, showing that energy has been developed ; there is a chem- 
ical action between the water and the oxide. We pour into sepa- 
rate vessels the liquids from the two jars, and to that from the 
phosphorus oxide we add some blue litmus solution, prepared by 
boiling litmus, a substance made from a peculiar moss, with water. 
The blue color instantly changes to red. We pour some of this 
red liquid into the water from the sodium experiment, and the 
blue color at once reappears. It is certain then that our two 
oxides have different properties, and the study of these and other 
oxides has shown that when oxygen combines with a non-metal- 
lic element, the resulting oxide usually combines with water, and 
forms a substance which changes blue litmus to red. Such sub- 
stances generally have a sour taste, and are called acids. On the 
contrary, the oxides of the metallic elements, if they have any 
effect, change the reddened litmus to blue, and are called basic 
oxides. Many oxides, however, have no effect on either red or 
blue litmus. 

62. When an oxide reacts with water, a body called a hydroxide 
or hydrate is formed. The acids containing oxygen are hydrates 
corresponding to non-metallic oxides, while the metallic oxides 



52 LESSONS IN CHEMISTRY. 

usually have corresponding metallic hydroxides, sometimes called 
metallic hydrates. We have seen the formation of sodium hydrox- 
ide by the action of sodium on water ; this same compound results 
from the action of water on sodium oxide, and we will notice that 
the only difference between the hydroxide and water is that the 
former contains an atom of sodium in place of one atom of 
hydrogen. 

Na 2 + H 2 = NaOH + NaOH 

Sodium oxide. Water. Sodium hydroxide. Sodium hydroxide. 

63. An analysis of the oxide of phosphorus which we have 
formed, shows that its molecule contains phosphorus combined 
with five atoms of oxygen ; it is therefore called phosphorus 
pentoxide* and, because the acid which it forms is called phos- 
phoric acid, the oxide is sometimes called phosphoric oxide. 
In general, the name of an oxygen compound is formed by putting 
oxide after the name of the other element, and to the word oxide 
is prefixed the Greek name of the number of atoms of oxygen in a 
molecule of the oxide. 

A monoxide contains one atom of oxygen, a dioxide contains 
two atoms of oxygen, a trioxide contains three, a tetroxide con- 
tains four, a pentoxide five. 

The word sesquioxide is sometimes used to indicate a com- 
pound whose molecule contains three atoms of oxygen and two 
atoms of the other element : sesqui means one and a half. Man- 
ganese sesquioxide contains Mn 2 3 . 

Sometimes an element forms more than one compound with 
oxygen. Nitrogen forms five ; and when we have learned that a 
molecule of each of these oxides contains two atoms of nitrogen, 
the names will at once indicate the composition of the molecules. 

Nitrogen monoxide, N 2 0. 
Nitrogen dioxide, N 2 2 (really NO). 
Nitrogen trioxide, N 2 3 . 
Nitrogen tetroxide, ~N 2 O i . 
Nitrogen pentoxide, N 2 5 . 

64. Frequently when there are only two oxides of an element, 
or when there are two of special importance, the word oxide is 

* Penta — five. 



NOMENCLATURE OF COMPOUNDS OF OXYGEN, ETC. 53 

not changed, but the name of the other element is made to end in 
ic or ous. There are only two oxides of mercury ; that containing 
the largest proportion of oxygen is called mercuric oxide, while 
that containing the least proportion is mercurous oxide. Their 
molecules contain 

Mercuric oxide, HgO. 
Mercurous oxide, Hg 2 0. 

Each of the two more important oxides of sulphur contains one 

atom of sulphur combined respectively with three and two atoms 

of oxygen. 

Sulphuric oxide, SO 3 . 
Sulphurous oxide, SO 2 . 

The oxide whose name ends in ic then contains a larger propor- 
tion of oxygen than that whose name ends in ous, and we should 
not use these terminations unless there be two oxides of the 
element. 

We can now understand what is meant when we say that water 
is hydrogen oxide, and we will presently learn the signification of 
all the names which we have been obliged to use. 

OZONE. 

65. Before, and sometimes during, a thunder-storm, there is 
often a peculiar odor in the air, and the same odor may be noticed 
near a good electric machine in operation. It has been found that 
the air has at the same time acquired more active oxidizing prop- 
erties than it had before. It will even bleach many coloring mat- 
ters. Part of the oxygen of the atmosphere has been 
changed to a body which we call ozone. 

66. We can produce this change by a simple experi- 
ment. Under the surface of some water we scrape the 
outside of a stick of phosphorus, so that it may be per- 
fectly free from oxide, and then put it into a bottle 
containing enough water to about half cover the phos- 
phorus, so that it may not take fire (Fig. 35). After 

it has stood for a little while, we dip into the air in the bottle a 
piece of paper that has been soaked in some starch boiled in water 
to which a little potassium iodide has been added. We see that 




54 



LESSONS IN CHEMISTRY. 



the paper at once becomes blue. Now let us put a drop of a solu- 
tion of iodine in alcohol on paper soaked in starch to which no 
potassium iodide was added. The same blue color appears. Po- 
tassium iodide is a compound of potassium and iodine : part of the 
oxygen of the air in the jar was converted into ozone; this took 
the potassium away from the iodine, and as soon as the latter be- 
came free it combined with the starch, producing the blue color. 
If we smell the air in the bottle, we find that it has a peculiar, 
and not very pleasant, odor, by which ozone may be identified as 
certainly as by the chemical test with potassium iodide and starch. 
It has been found that these same phenomena are produced 
by pure oxygen gas through which electrical sparks have been 
passed (Fig. 36), and that by the passage of such sparks the vol- 
ume of the oxygen is diminished, 
while its weight of course does not 
change. The increase in density so 
observed has shown that ozone is 
half again as heavy as oxygen : when 
ozone is heated, it is converted into 
ordinary oxygen, and the volume is 
expanded in the same proportion. 
Chemists have consequently been led 
to believe that while ordinary oxygen 
contains two atoms in its molecule, a 
molecule of ozone contains three such 
atoms. We may consider, therefore, 
that if a molecule of ordinary oxygen 
is represented by the formula O 2 , O 3 
represents a molecule of ozone. 

67. Let us see why ozone possesses 
more active powers than oxygen. 
When we pass electric sparks through oxygen, we decompose its 
molecules, and the energy of electricity is transferred to the atoms, 
which it enables to combine by threes, instead of by twos. Phos- 
phorus is gradually oxidized by oxygen, but one atom of phos- 
phorus does not combine with whole molecules of oxygen: we 




Fig. 36. 



HYDROGEN DIOXIDE. 55 

shall in time learn that in this case two atoms of phosphorus take 
three atoms of oxygen ; that would be a molecule and a half; but, 
while the energy developed by the rapid combustion of phosphorus 
appears as heat and light, the energy developed by the slow com- 
bustion of the phosphorus is transferred to the odd atom of oxy- 
gen, and enables it to combine with two other atoms set free from 
molecules in the same manner. 

6P + 60 2 = 3P 2 3 + O 3 

Phosphorus, Oxygen, Phosphorus trioxide, Ozone, 

six atoms. six molecules. three molecules. one molecule. 

As ozone contains more energy than oxygen, its properties are 
more energetic ; thus, a bright silver coin suspended in ozone 
soon becomes tarnished by oxidation. 

We shall have occasion to study many actions of this kind, where 
the energy evolved by the combination of certain atoms is trans- 
ferred to other atoms, giving them more active properties than 
they had before. 

When ozone oxidizes other bodies, in most cases only one of its 
atoms is used in the oxidation ; the other two unite to form a 
molecule of oxygen. 

When the moist potassium iodide was decomposed by ozone, 
potassium hydroxide was formed ; its molecule contains KOH, 
and we see that the water present must have taken part in the 
reaction, which we may write 

2KI + H20 + O 3 = 2KOH + O 2 + I 2 

Potassium iodide. Potassium hydroxide. 

68. Ozone is frequently produced in slow combustions. By 
great cold it may be converted into a sky-blue liquid. It 
is destroyed, that is, converted into oxygen, by a temperature 
of 290°. Its oxidizing powers are sometimes employed for bleach- 
ing and disinfecting, the ozone in these cases being produced by 
electricity. 

HYDKOGEN DIOXIDE, H 2 2 . 

69. We introduce some pulverized barium dioxide, a compound 
whose molecule contains one atom of the metal barium and two 
atoms of oxygen, into a small flask containing some cold dilute hy- 
drochloric acid ; as the solid dissolves, a solution of barium chloride 



56 LESSONS IN CHEMISTRY. 

is formed, while the hydrogen of the hydrochloric acid and the 

oxygen of the barium dioxide combine to form a compound called 

hydrogen dioxide, which remains dissolved in the liquid. 

BaO 2 + 2HC1 ,= BaCl 2 + H 2 2 

Barium dioxide. Hydrochloric acid. Barium chloride. Hydrogen dioxide. 

The separation of the hydrogen dioxide from the barium chlo- 
ride is not an easy matter, but the latter compound will not inter- 
fere with our experiments. 

We pour some of the liquid on a little manganese dioxide ; at 
once a brisk effervescence takes place, and by the aid of a match- 
stick bearing a spark, we are shown that the tube is filled with 
oxygen. The hydrogen dioxide has been decomposed into water 
and oxygen ; but the manganese dioxide remains unchanged ; it is 
probably in fact converted into a higher oxide, but the addi- 
tional oxygen is at once taken away from this oxide by another 
atom of oxygen with which it forms a molecule of the gas. 

In another tube, we pour a little of our solution on some black 

lead sulphide : the color quickly changes to white, and no gas is 

given off, for the lead sulphide is converted into lead sulphate, 

while water is formed. 

PbS + 4H20 2 = PbSO 4 + 4H 2 

Lead sulphide. Lead sulphate. 

We now pour a little hydrogen dioxide into some purple solution 
of potassium permanganate which has been acidified with sulphu- 
ric acid. The liquid becomes colorless ; at the same time bubbles 
of oxygen are disengaged, and may be identified by the usual test. 
In this case an atom of oxygen, very loosely held by the other 
atoms in the hydrogen dioxide, has combined with another atom 
from the potassium permanganate, which is very rich in oxygen, 
and a molecule of free oxygen is given off, while water is formed 
as before. 

We mix some of the hydrogen dioxide liquid with a little 
yellow solution of potassium dichromate ; we then quickly pour 
in a quantity of ether, and briskly shake the tube ; the ether 
being lighter than the water, comes to the top of the latter, in 
which it is almost insoluble, and this layer of ether has a dark 



CHLORINE. 57 

blue color. It contains perchromic acid, a body which is formed 
by the oxidation of the potassium dichromate ; but with hydrogen 
dioxide this perchromic acid behaves just like potassium perman- 
ganate ; unless it is at once removed from the liquid in which it is 
formed, its oxygen is taken away, and a green liquid containing a 
lower oxide of chromium is obtained. 

70. We may then conclude that hydrogen dioxide acts in three 
ways with other substances : sometimes it is reduced, that is, part 
of its oxygen is taken away, while the other body remains un- 
changed, as is the case with manganese dioxide ; sometimes the 
second substance is oxidized, as with the lead sulphide and potas- 
sium dichromate ; sometimes both the hydrogen dioxide and the 
other body are deoxidized, as in the cases of potassium perman- 
ganate and perchromic acid. 

Pure hydrogen dioxide is a syrupy, colorless liquid, without 
odor, and having a density of 1.45. It is slowly decomposed into 
water and oxygen at ordinary temperatures, with brisk efferves- 
cence at 100°, and explosively if dropped on a surface heated to 
higher temperatures; in vacuo it may be distilled without de- 
composition. Owing to the readiness with which it gives up oxy- 
gen, hydrogen dioxide will destroy organic coloring matters and 
germs of disease. It is extensively manufactured for bleaching 
silk, ostrich feathers, and human hair ; an aqueous solution con- 
taining three per cent, is used in medicine as an antiseptic. 

Hydrogen dioxide and ozone undergo mutual decomposition, 

vvater and free oxygen being formed. 

H20 2 + O 3 - H*0 + 20 2 

Hydrogen dioxide. Ozone. Water. Two molecules of oxygen. 



LESSON VIII. 

CHLORINE— CHLORIDES. 

Atomic weight, 35.5. Symbol, CI. 

71. The element chlorine has such strong affinities for other 
elements that it does not exist free in nature, but is always found 



58 



LESSONS IN CHEMISTRY. 



in combination. The most important of its compounds is common 
salt, which contains sodium and chlorine, and of which enormous 
quantities exist in the ocean, and in salt springs and salt mines. 
We do not usually prepare chlorine directly from salt, but from 
hydrochloric acid, the latter being prepared from the salt itself. 

We mix in a glass flask some strong hydrochloric acid with 
about one-sixth its weight of manganese dioxide, and, after adapt- 
ing to the flask a cork through which passes a tube for the exit 
of the gas, and another tube called a safety-tube, we gently heat 
the mixture over a flame (Fig. 37). The safety-tube (A), which 




Fig. 37. 



is bent around on itself and has a little bulb blown on the bend, 
enables us to add more acid if necessary, and at the same time if 
there should be too much pressure in the flask it allows the gas 
to escape through the little acid which we must pour into it: 
on the contrary, when the flask cools and the gas contracts in vol- 
ume, air may enter through the safety-tube, and any liquid into 
which we may wish the end of the delivery-tube to dip, will not 
be drawn back into the flask. We may dry our chlorine gas by 
passing it through a bottle containing either calcium chloride or 
strong sulphuric acid, or we may pass it directly into a bottle. 



CHLORINE. 59 

Chlorine dissolves in water, and we collect it by downward dry 
displacement ; for it is a heavy gas, and when we pass the tube 
through which it flows to the bottom of a jar, the chlorine grad- 
ually forces the air out at the top. We might collect it over salt 
water in the pneumatic trough, as it does not dissolve in salt 
water. We can easily see when the jar is full, for the gas has a 
greenish-yellow color. While we are filling several jars, which 
we cover with glass plates as soon as they are filled, we may ex- 
amine the chemical change by which chlorine is formed. Man- 
ganese dioxide contains two atoms of oxygen, and this oxygen 
combines with the hydrogen of the hydrochloric acid, forming 
water. Two atoms of oxygen require four atoms of hydrogen, 
and for these four atoms we will need four molecules of hydro- 
chloric acid, each of which contains one atom of chlorine and one 
of hydrogen. The atom of manganese combines with two atoms 
of chlorine, forming a body called manganese chloride, and as 
there were four chlorine atoms in the hydrochloric acid, two of 
these will pass off as gas. We may write the reaction, 

MnO 2 + 4HC1 = MnCl 2 + 2H20 + CI 2 

Manganese dioxide. Hydrochloric acid. Manganese chloride. Chlorine. 

72. Properties. — Chlorine is a greenish-yellow gas, having an 
unpleasant, suffocating odor. We must be careful not to breathe 
it in a too undiluted form, for it causes violent coughing, and irri- 
tates the lungs. It is 2.45 times as heavy as an equal bulk of 
air, or 35.5 times as heavy as an equal volume of hydrogen. Its 
atomic weight is also 35.5 : there are many other elements whose 
atomic weights and densities (when in the form of gas) compared 
to hydrogen are the same, and we must suppose that the mole- 
cules of such elements are like those of hydrogen in that each 
contains two atoms (§ 46). Chlorine dissolves in water : at ordi- 
nary temperatures, one litre of water will dissolve about two 
and a half litres of the gas. It may be liquefied at 15° by a 
pressure of four atmospheres, that is, four times as great as the 
ordinary pressure of the air. 

Chlorine possesses very great affinity for the other elements, 
and the compounds which it forms with them are called chlorides. 



60 



LESSONS IN CHEMISTRY. 



Over one of our jars of the gas we place a piece of coarse wire 
gauze, through which we sprinkle some finely-powdered antimony. 
Each little particle burns, and we have a shower of fire, while a 
heavy cloud of white smoke settles in the jar: this smoke is 
antimony chloride. Into another jar we throw some pieces of 
Dutch leaf, a very thin brass used for cheap gilding : this also 
burns, and the copper and zinc of which the Dutch metal was 
composed are converted into copper chloride and zinc chloride. 
We may burn some thin copper in the same manner. We put a 
small piece of phosphorus in a deflagrating-spoon, and lower this 
into another jar : it burns with a pale flame into phosphorus 
chloride. 

73. Of all the elements, hydrogen is that for which chlorine 
possesses the most remarkable affinities. In a room lighted only 
by a candle, we have mixed over salt water equal volumes of 
chlorine and hydrogen, and, after drying this mixture by passing 
it through a tube containing pumice-stone and sulphuric acid, we 
have filled with it some little bulbs, blown on thin glass tubes, 
and then sealed the ends of the tubes with little plugs of paraffin. 
It is easy to fill the bulb ; we connect one end of it by a rubber 

tube on which is a pinch 
(A), to the tube of the 
bell-jar in which the mix- 
ture is made ; then on 
pressing the jar into the 
salt water and loosing the 
pinch, the gas is forced 
through the bulb (Fig. 
Fig. 38. 38). We keep these bulbs 

carefully covered from the 
light. We now uncover one, put it behind a sheet of glass, and 
then, standing at a little distance, burn a piece of magnesium wire. 
Instantly there is an explosion ; the hydrogen and chlorine have 
combined. The combination is brought about in the same manner 
by direct sunlight, and more gradually by diffuse daylight. 

Chlorine does not support ordinary combustion, for it does not 




CHLORINE. — CHLORIDES. 61 

combine directly with carbon, and ordinary combustibles contain 
hydrogen and carbon ; but their hydrogen may burn in chlorine. 
We put a lighted taper into a jar of chlorine, and the flame be- 
comes red and smoky : the chlorine combines with the hydrogen 
of the wax, but the carbon separates in the form of smoke. 

Chlorine even decomposes many compounds containing hydro- 
gen, taking away that element to form hydrochloric acid. The 
solution of chlorine in water is decomposed by sunlight, the oxy- 
gen being set free. 

2C1 2 + 2H 2 = 4HC1 + O 2 

Into a jar with straight sides, filled with chlorine, we rapidly 
introduce a paper saturated with turpentine, and quickly replace 
the cover of the jar. There is a flash of red light, and a cloud of 
smoke. Turpentine is a compound of carbon and hydrogen only : 
the chlorine combines with the hydrogen, and the carbon forms 
the smoke. We find after the experiment that the paper is not 
burned : we use a plain straight jar, because it is easily cleaned by 
rubbing with a little turpentine. 

We pour some blue litmus-water into a jar of chlorine ; the 
blue liquid becomes colorless. In another jar we suspend a piece 
of moist colored calico, and it quickly fades. Chlorine possesses 
bleaching properties, and these properties are due to the decompo- 
sition of the dye-stuffs, nearly all of which contain hydrogen that 
the chlorine may remove. For the same reason chlorine is a val- 
uable disinfectant, for most unpleasant and unwholesome odorous 
matters are compounds of hydrogen. 

74. Chlorides. — The binary compounds of chlorine are called 
chlorides, and the same prefixes which are used for the names of 
the oxides are employed also to indicate the number of chlorine 
atoms in a molecule of the compound ; thus, phosphorus trichloride 
contains PCI 3 . In general, these prefixes are used to indicate the 
number of atoms of the second named element with which one or 
more atoms of that first named are combined. When there are 
only two chlorides which are important, the terminations ous and 
ic designate which contains the least and the greatest proportion 
of chlorine (see § 64). Mercurous chloride is Hg 2 CP ; mercuric 



62 LESSONS IN CHEMISTRY. 

chloride is HgCP ; this nomenclature also is of general applica- 
tion. 

The metallic chlorides are all soluble in water, with the excep- 
tion of silver chloride, AgCl, mercurous chloride, Hg 2 Cl 2 , and 
cuprous chloride, Cu 2 CP. Lead chloride is only slightly soluble. 
The non-metalic chlorides, as a rule, are decomposed by water, 
and in such cases part or all the chlorine combines with hydro- 
gen, forming hydrochloric acid. Thus, phosphorous chloride, 
PCI 3 , yields phosphorous acid and hydrochloric acid. 

PCI 3 + 3H 2 = H3P0 3 + 3HC1 

Phosphorous chloride. Phosphorous acid. Hydrochloric acid. 

75. We pour a few drops of a solution of silver nitrate in pure 

water into some water in which common salt, which is sodium 

chloride, has been dissolved. At once a white precipitate forms, 

for, while sodium nitrate now exists in solution, silver chloride is 

formed, and this is insoluble. 

AgNO 3 + NaCl = NaNO 3 + AgCl 

Silver nitrate. Sodium chloride. Sodium nitrate. Silver chloride. 

The precipitate darkens on exposure to light, and this reaction 
enables us to determine whether a body contains or does not con- 
tain a chloride. All solutions of chlorides give the white precipi- 
tate, which, we may add, dissolves if we pour off most of the 
liquid and then shake the white powder with strong ammonia- 
water. 



LESSON IX. 

HYDROCHLORIC ACID.— ACIDS.— SALTS. 

76. Hydrochloric Acid, HC1. — We have seen that this com- 
pound is formed by the direct union of chlorine and hydrogen, 
and by the action of water on certain chlorides. Many chlorides 
are decomposed by water at high temperatures, and in this manner 
some mineral chlorides existing in the rocks cause hydrochloric 



HYDROCHLORIC ACID. 



63 



acid to be formed in certain volcanic regions, where it mixes with 
the other gases that are emitted. 

77. Preparation. — Hydrochloric acid is made by the action of 
sulphuric acid on common salt, the sodium of the salt changing 
places with the hydrogen of the sulphuric acid. 

We put some pieces of rock-salt in a flask like that in which 
we made chlorine, and, if we wish a solution of the gas. we con- 
nect our delivery-tubes with a series of bottles containing water, 
through which the gas will be forced to bubble (Fig. 39). If we 




Fig. 39. 

wish the dry gas, we dry it as we did the chlorine. When all is 
ready, we pour through the safety-tube sulphuric acid which we 
have previously diluted with an equal volume of water, and im- 
mediately the gas begins to come off. When the reaction becomes 
tranquil, we must heat the mixture. 

One molecule of sulphuric acid contains two atoms of hydrogen, 
and may be made to yield one or two molecules of hydrochloric 
acid, by reacting with one or with two molecules of salt. 

NaCl + H2SO* = HC1 + NaBSO 4 

Sodium chloride. Sulphuric acid. Hydrochloric acid. Sodium acid sulphate. 

2NaCl + H 2 SO± = 2HC1 + Na«SO* 

Sodium sulphate. 



64 LESSONS IN CHEMISTRY. 

This reaction is operated on an enormous scale in Europe, 
where the sodium sulphate is afterwards converted into soda or 
sodium carbonate. 

78. Properties. — Hydrochloric acid is a colorless, pungent, and 
suffocating gas. Its density compared to hydrogen is 18.33, suf- 
ficiently near that which would be indicated by half its molecular 
weight, which is 36.5 (see § 48). It is very soluble in water, 
and if under the surface of water we remove the cork from a bot- 
tle filled with the gas, the water at once rises and fills the bottle. 
At 0° one litre of water will dissolve 500 litres of hydrochloric 
acid. The strongest hydrochloric acid of commerce, commonly 
called muriatic acid, contains about 34 per cent, of the gas. Like 
the gas, it produces fumes in the air, some gas escaping from it 
and condensing the moisture in the atmosphere. 

Hydrochloric acid is a strong acid. A drop or two of the solu- 
tion will redden a large quantity of blue litmus. We slowly pour 
some hydrochloric acid into a strong solution of sodium hydrox- 
ide : a white powder soon separates, and we can satisfy ourselves 
by tasting it that this is common salt. Water also is formed. 
NaOH + HC1 = IPO + NaCl 
Sodium hydroxide. Sodium chloride. 

With oxides of the metals, hydrochloric acid acts in the same 
manner, water and a chloride being formed. 

HgO + 2HC1 = HgCl 2 + H20 

Mercuric oxide. Mercuric chloride. 

We have seen that zinc liberates the hydrogen from hydro- 
chloric acid : many other metals act likewise. 

ACIDS AND SALTS. 

79. An acid is a compound containing hydrogen which is ca- 
pable of being replaced by a metal, forming a body which is called 
a salt. Although salts may be formed in various manners, we have 
an exact definition : a salt represents an acid whose hydrogen has 
been partly or wholly replaced by metal. Hydrochloric acid is an 
example of a binary acid, but the few binary acids which we shall 
study have not all as energetic properties as hydrochloric acid. 
The salts formed by hydrochloric acid are of course the chlorides. 



HYPOCHLOROUS OXIDE AND ACID. 65 

80. Hypochlorous Oxide and Acid. — When chlorine is passed 
over cooled mercuric oxide, mercuric chloride is formed, and the 
oxygen which separates from the mercury combines with chlo- 
rine, forming a gas which may be condensed to a yellow liquid 
by passing it into a bottle surrounded by a freezing mixture of 
ice and salt. 

HgO + 2C1 2 = HgCl 2 + C1 2 

Mercuric oxide. Mercuric chloride. Hypochlorous oxide. 

This is hypochlorous oxide ; it is a dangerous body, and often 
explodes without warning. 

It reacts with water in a manner which we must study. A 
molecule of the oxide and a molecule of water interchange a chlo- 
rine atom for a hydrogen atom, and a compound called hypochlo- 
rous acid is formed. 

ClOCl + HOH = HOC1 ClOH 

Hypochlorous oxide. Water. Hypochlorous acid. Hypochlorous acid. 

This is an oxygen acid, and we may consider that it is com- 
posed of an atom of chlorine combined with the residue of a mol- 
ecule of water from which one atom of hydrogen has been removed. 
This residue would be OH, and, because the atom of oxygen has 
not enough hydrogen to satisfy the affinities, it is not a molecule ; it 
cannot exist except as part of a molecule ; that is, combined with 
some other atom. It is called, for convenience' sake, hydroxyl, and 
all oxygen acids contain this group of two atoms, hydroxyl. In- 
deed, all the compounds we call hydroxides contain the group 
hydroxyl ; thus, potassium hydroxide is KOH. 

81. We have had occasion to notice the names hydrochloric 
acid, hypochlorous acid, sulphuric acid. We have seen that hy- 
drochloric acid produces binary salts : the names of binary com- 
pounds, with the exception of acids, end in ide, and we can now 
even understand that a sulphide is a compound containing sulphur 
and one other element. But hypochlorous acid and sulphuric 
acid are not binary compounds ; they may be formed respectively 
by the action of hypochlorous and sulphuric oxides on water. The 
first of these actions we have studied : the second we may write 

SO* + H 2 = H 2 SO*. 
5 



66 LESSONS IN CHEMISTRY. 

When the hydrogen of either of these oxygen acids is replaced 
by metal, how shall we name the resulting salts ? Chemists have 
agreed that the termination ic shall be changed to ate, and ous 
shall be changed to ite. This is a simple nomenclature. The salts 
of sulphuric acid must be sulphates ; those of nitric acid, nitrates ; 
those of permanganic acid, permanganates ; those of hypochlorous 
acid, hypochlorites ; those of sulphurous acid, sulphites. We see 
also that the chlorates must be the salts of chloric acid ; the ar- 
senites, those of arsenious acid. 

82. Hypochlorites. — Solutions of the hypochlorites of potas- 
sium and sodium are useful as disinfecting and bleaching liquids. 
They are made by passing chlorine gas into a rather dilute solu- 
tion of potassium hydroxide or sodium hydroxide ; at the same 
time water is formed, and a chloride, which remains in solution. 

2NaOH + CI 2 = NaOCl + NaCl + H 2 

Sodium hydroxide. Sodium hypochlorite. Sodium chloride. 

Such a liquid quickly removes the stains of wine and fruits 
from linen, and also deodorizes offensive matters. 

Bleaching powder, or chlorinated lime, is made by passing chlo- 
rine gas over slaked lime. Its solutions contain calcium hypo- 
chlorite, Ca(ClO) 2 , and may be substituted for the liquids which 
we have just mentioned. The bleaching and disinfecting by these 
substances are due to their decomposition, which we may suppose 
first sets free hypochlorous acid, and this attacks the coloring 
matter or offensive substance, removing hydrogen ; the chlorine 
atom will take one atom of hydrogen, forming hydrochloric acid, 
and the group OH takes another atom, forming water. We may 
understand this by examining the reaction between hydrochloric 
and hypochlorous acids, which yields chlorine and water. 
HCIO + HCl = H*0 + CI 2 

83. Chlorates. — We pass a current of chlorine gas into a 
strong solution of potassium hydroxide, and a white solid matter 
soon appears in the liquid ; when this no longer increases in bulk, 
we stop the chlorine, heat the liquid until it boils, and if all of the 
solid dissolves, we evaporate it until a considerable quantity of 
this matter again separates. We now allow it to settle a moment, 



CHLORIC ACID. — CHLORATES. 67 

and pour off the clear liquid : as this cools, shining little crystals 

separate in rhomboidal plates. These are potassium chlorate, and 

we have been obliged to separate them from potassium chloride, 

which, together with water, is also formed during the experiment. 

6KOH + 3C1 2 = 5KC1 + KCIO 3 + 3H 2 

Potassium hydroxide. Potassium chloride. Potassium chlorate. 

Potassium chlorate is the most important salt of chloric acid, 
HCIO 3 , which we might prepare from the salt by a troublesome 
process. Potassium chlorate is not very soluble in cold water, 
but is very soluble in boiling water. Its solution is excellent as a 
gargle for sore throat, but must not be swallowed, for it is poison- 
ous. We have seen that potassium chlorate is decomposed by 
heat, yielding oxygen : it readily gives up its oxygen, for the 
chlorine has a much stronger affinity for the potassium than for 
the oxygen, which appears to hold the chlorine and potassium 
atoms together. 

84. Chloric acid, which would be set free by the action of 
stronger acids on potassium chlorate, is at once decomposed under 
such circumstances if oxidizable substances be present. On a 
mixture of equal parts of potassium chlorate and sugar, powdered 
separately, we let fall a drop of strong sulphuric acid. The mix- 
ture at once takes fire and burns vividly, the potassium chlorate 
furnishing the oxygen for the combustion of the 
sugar. 

Into a tall jar, filled with water, we throw some 
crystals of potassium chlorate, and on them a small 
piece of phosphorus ; then, by means of a funnel- 
tube which passes to the bottom of the jar, we pour 
some strong sulphuric acid on the chlorate. The 
chloric acid set free is decomposed by the phos- 
phorus and causes its combustion under the water 
(Fig. 40). 

We put into a mortar a piece of sulphur about as 
large as a match-head, and a crystal of potassium 
chlorate of the same size ; then we rub them briskly together, being 
careful to keep the mortar far enough from the face, and soon there 




68 LESSONS IN CHEMISTRY. 

is a loud report ; the sulphur has been oxidized and the potassium 
chlorate decomposed. If we used larger quantities of these sub- 
stances in our experiment, we might break the mortar, and possibly 
injure our person. 



LESSON X. 
BROMINE.— IODINE.— FLUORINE. 

85. Bromine, Br = 80. — In a long tube closed at one end, 
we dissolve in a little water a few crystals of a white substance, 
called potassium bromide, and then pour in some chlorine-water, 
which we have prepared by passing chlorine through water con- 
tained in bottles such as were used in the preparation of the 
solution of hydrochloric acid : we now add a considerable propor- 
tion of ether, and shake the tube after closing it with the finger. 
The liquid becomes brown, and after standing a few minutes, the 
ether, which is not very soluble in water, comes to the surface, 
and its color is red, while the water has become colorless. The 
potassium bromide, a compound of potassium and bromine, has 
been decomposed by the chlorine, and potassium chloride formed 
in the solution, while the bromine set free has been dissolved by 
the ether, in which it is much more soluble than in water. We 
may write the reaction, 

2KBr + CI 2 = 2KC1 + Br 2 

Potassium bromide. Potassium chloride. Bromine. 

86. The compounds of bromine with potassium, sodium, and 
magnesium, which compounds are called bromides of those metals, 
are found in the waters of many salt springs, and exist in small 
quantity in the water of the ocean. As they are much more 
soluble in water than common salt, they remain dissolved when 
most of the salt has been separated by evaporating the liquid, and 
from their concentrated solution so obtained the bromine is sepa- 
rated by heating the liquid with sulphuric acid and manganese 



BROMINE. IODINE. 69 

dioxide. Supposing all of the bromine to exist as potassium 
bromide, manganese sulphate, potassium sulphate, and water are 
formed ; the bromine distils, and is condensed in suitable apparatus. 

2KBr + MnO 2 + 2H 2 S0 4 = R2SO* + MnSO* + 2H 2 + Br 2 
Potassium Manganese Sulphuric Potassium Manganese 
bromide. dioxide. acid. sulphate. sulphate. 

Chlorine is similarly separated from chlorides by heating them 
with manganese dioxide and sulphuric acid. 

87. Bromine is a dark-red liquid, having an exceedingly irri- 
tating odor. Its density is 3.2. It freezes at — 7.3°, and boils at 
59° ; it is very volatile at ordinary temperatures. It dissolves in 
about thirty times its weight of water at 15°, and is quite soluble 
in ether, chloroform, and carbon disulphide, liquids which dissolve 
many substances that are not soluble in water. 

Bromine closely resembles chlorine in its chemical reactions, 
but its affinities are not as powerful. Its solution in water will 
bleach litmus, and other coloring matters, but more feebly than 
chlorine. We pour a little bromine into a deep test-tube and 
drop in a piece of warm copper foil which is instantly converted 
into copper bromide with the production of heat and light. 

Bromine combines with hydrogen, forming hydrobromic acid, 
HBr, a gas which closely resembles hydrochloric acid. 

Bromine is exceedingly corrosive to animal tissues, and is 
sometimes employed as a caustic in surgery. It disinfects like 
chlorine, for which the dilute aqueous solution of bromine is 
often substituted. 

The atomic weight of bromine is 80, and the density of its 
vapor compared to hydrogen is also 80, showing that a molecule 
of bromine contains two atoms. 

88. Iodine, I = 127. — In a tube like that which we used in 
the experiment with potassium bromide, we dissolve a little potas- 
sium iodide in water, add chlorine-water as before, and then, in- 
stead of ether, we pour in some carbon disulphide. After shaking 
the tube, and allowing it to stand, the carbon disulphide, being 
heavier than the water, is found at the bottom with a beautiful 
purple color. Were we to pour off the watery liquid and allow 



70 LESSONS IN CHEMISTRY. 

this carbon disulphide to evaporate in a shallow dish, it would leave 
a brownish-gray matter, which is iodine, and which the chlorine 
has driven out of the potassium iodide, just as it separated the 
bromine from the potassium bromide. 

89. Like bromine, iodine is found combined with potasssium, 
sodium, and magnesium in the waters of some springs, and in sea- 
water. Compounds of iodine also exist in the sodium nitrate 
found in large deposits in Chili, and, being very soluble in water, 
remain in the mother-liquor, as it is called, from which this 
sodium nitrate has been crystallized for its purification. Iodine is 
obtained from these liquids, and from the ashes of sea-weeds ; the 
sea-weeds are burned, and the iodides which are dissolved out of 
the ashes by water, are treated just as the bromides are treated 
for the preparation of bromine. Iodine may also be separated by 
adding nitric acid to the solution of an iodide, and we may make 
the experiment by pouring a little nitric acid on some potassium 
iodide solution in a test-tube. Potassium nitrate is formed, and 
iodine deposits as a dark powder : the red vapors that are given 
off are a compound of nitrogen and oxygen, which we will study 
in good time. 

90. Iodine is purified by sublimation ; that is, heating it, and 
condensing the vapor. When pure, it is iu crystalline, bluish- 
gray plates, much like scales of plumbago. Its density is 4.95. 
It melts at 114°, and boils at 184°. We carefully heat a few 
small scales of iodine in a large glass flask, which soon becomes 
filled with a magnificent purple vapor. This vapor is so heavy 
that we may pour it out on a piece of cold glass, where it con- 
denses in minute crystals. Below 600° the density of iodine 
vapor is 127, but above 1400° only about half as much; each 
I 2 molecule breaks up into two single atoms. 

Iodine is very slightly soluble in water : one part of iodine re- 
quires 7000 parts of water to dissolve it, and yet the solution has 
a distinct brown color. It dissolves readily in alcohol, ether, 
chloroform, and carbon disulphide, and the color of the solution 
depends on the solvent ; that in alcohol is brown, but that in 
chloroform is violet. 



FLUORINE. 71 

91. We have made some thin starch paste by boiling starch 
with water, and we pour some of this into two test-tubes : to the 
first we add a few drops of a solution of iodine in water, and the 
liquid becomes dark blue ; to the other we add a drop or two of a 
solution of potassium iodide, and no color is produced. Starch is 
dyed a blue color by iodine, but the iodine must be free ; on add- 
ing a few drops of chlorine-water or nitric acid to the second 
tube the potassium is removed from the iodine, and the blue 
color at once appears. This is the test for iodine. 

Iodine combines with hydrogen to form hydriodic acid, HI, a 
gas whose properties are much like those of hydrochloric acid. It 
is made by heating water with iodine and amorphous phos- 
phorus. 

92. Analogies of CI, Br, and I. — On comparing the three elements which 
we have just considered, we find that while one is a gas, another liquid, and 
the third solid, still the corresponding compounds formed by each are much 
alike in chemical nature; that is, the composition and reactions of the mole- 
cule. The compounds with hydrogen each contain one atom of hydrogen com- 
bined with one of the other element : if the power to combine with one atom 
of hydrogen be taken as the measure of the combining power of any atom, 
the atoms of chlorine, bromine, and iodine must have equal powers. Since 
an atom of each of these elements combines with only one atom of hydrogen, 
they are said to be monatomic elements in their compounds with hydrogen. 
But their affinities, or energies of combination, for hydrogen are not alike : 
chlorine will take the hydrogen away from hydrobromic acid, and bromine 
will take it away from hydriodic acid. This is also the order of their affinity 
for the metals, but in the number of atoms of either chlorine, bromine, or 
iodine which will combine with one atom of another element, the three are 
exactly alike. 

In this last respect the next element resembles the three which we have just 
studied. 

93. Fluorine, F = 19. — We have coated one side of a glass 
plate with wax, and in the wax we trace a design with a sharp 
point, taking care that our lines go quite through to the glass. 
In a shallow dish made of sheet lead, we mix, by the aid of a 
wooden stick, some powdered fluor-spar, which is a mineral, with 
strong sulphuric acid ; over this we place our glass containing the 
design, with the waxed side down, and we gently warm the dish 
(Fig. 41). In a few minutes we remove the glass, and, after 



72 



LESSONS IN CHEMISTRY. 



gently warming it, rub off the wax : we find that the design is 
permanently etched into the glass. The fluor-spar is a compound 
of the elements fluorine and calcium, and the sulphuric acid has 




Fig. 41. 

decomposed it, forming a vapor called hydrofluoric acid, a com- 
pound of hydrogen and fluorine. 

CaF2 + H2S0 4 = CaSO* + 2HF 

Calcium fluoride. Sulphuric acid. Calcium sulphate. Hydrofluoric acid. 

Hydrofluoric acid may be condensed to a liquid, and it may 
be dissolved in water, but neither the liquid nor the solution can 
be kept in glass bottles, because fluorine has an extraordinary 
affinity for the silicon which forms part of the glass, and would 
combine with that element, destroying both bottle and acid. It 
is to this affinity that we owe the etching of our glass plate. 
Bottles of india-rubber or of lead are used to contain hydrofluoric 
acid, for it does not attack those substances. The graduations on 
delicate chemical apparatus, such as the eudiometers we have seen, 
are etched into the glass by this acid. Hydrofluoric acid is very 
corrosive, and we must be careful in its use. 

The powerful affinities of fluorine long prevented its isolation. It has been 
obtained by electrolyzing anhydrous hydrofluoric acid at low temperatures in 
vessels of platinum. It is a yellowish gas possessing a penetrating odor. It 
liquefies at very low temperatures. The density is 18.2. 

With hydrogen it combines with explosive violence even in the dark and at 
low temperatures. It decomposes water with formation of hydrofluoric acid 
and ozone, and displaces chlorine from its compounds. Most of the non- 
metallic elements and nearly all the metals burn in fluorine : even gold and 
platinum combine with it at higher temperatures, but for oxygen it seems to 
have no affinity. 

Besides fluor-spar, there is another important compound of fluorine found 
in nature ; it is the mineral cryolite, which is a compound of sodium fluoride 
with aluminium fluoride, 3NaFAlF 3 . 



SULPHUR. — HYDROGEN SULPHIDE. 



73 



LESSON XL 



SULPHUR.— HYDROGEN SULPHIDE. 



94. Sulphur, S = 32. — We are all familiar with sulphur, or 
brimstone. In some localities it is found pure or very impure 
and mixed with the soil : especially is this the case in volcanic 
countries. Besides this free or native sulphur, as it is called, sul- 




phur is found combined with many metals, and the compounds 
are called sulphides. 

Crude sulphur comes in large quantities from Sicily, where it 
is obtained by heating the ore so that the sulphur melts and 
flows from the earthy matters with which it is mixed. It is 
refined by distilling it in an apparatus consisting of an iron 
boiler (A, Fig. 42), above which is a reservoir (C) where the 
sulphur is first melted by the waste heat, and from which 
it runs into the boiler. The sulphur vapor enters a large 



74 LESSONS IN CHEMISTRY. 

chamber (B), and after condensing runs down on the floor, which 
is inclined so that the melted sulphur may be drawn off at a tap 
(H). While the walls of this chamber are yet cold, the sulphur 
condenses in a fine yellow powder, which is sold as flowers of sul- 
phur ; but when the chamber becomes heated, the condensed sul- 
phur melts, and after being drawn from the opening is cast in 
cylindrical moulds, in which it solidifies and becomes roll sulphur. 

Large quantities of sulphur are also obtained by distilling iron 
pyrites, a compound which contains iron and sulphur, and which 
gives up part of its sulphur when it is heated. 

95. Properties. — Sulphur is a brittle, lemon-yellow solid, 
having neither taste nor odor. It is a bad conductor of electricity 
and heat : a piece of roll sulphur held firmly in the hand produces 
a curious crackling noise, because the outside becomes warm, and 
its expansion causes it to crack before the heat can be conducted 
to the interior. Sulphur is not soluble in water, is very slightly 
soluble in alcohol and ether, but dissolves readily in carbon disul- 
phide. When heated, it melts at 114.5°, and becomes a mobile, 
amber-colored, and transparent liquid. 

We melt some sulphur in an earthen crucible, and, as soon as it 
has all melted, we allow it to cool until a crust forms over the 
surface. We now make a hole in the crust, and pour out the sul- 
phur which has not solidified. On breaking off the crust, we find 
the interior of the crucible lined with beautiful, transparent crys- 
tals, which on close examination we determine to be monoclinic 
prisms. In a glass flask we melt some more sulphur, but after 
it has melted we keep on heating it : we see that the color 
becomes darker, and the liquid thicker. When its temperature 
reaches 220°, we can turn the flask upside down and the sulphur 
scarcely runs on the sides. At about 260° it again becomes 
liquid, and at 448° it boils, forming a red vapor. We now 
pour it into cold water, moving the flask so that all does not 
fall in the same place. On taking the sulphur from the water, 
we find that its properties are much changed : it is transparent 
and very elastic ; we pull it out in long threads. This curious 
form, which is called soft sulphur, is due to a molecular condition 



SULPHIDES. — HYDROGEN SULPHIDE. 75 

of the element ; we must believe that its molecules contain more 
energy than those of ordinary sulphur, for if we gradually heat it, 
it at once becomes opaque and brittle, and at the same time much 
hotter than we have heated it. It changes spontaneously in this 
manner after we have kept it a few hours. Soft sulphur is 
amorphous ; that is, has no crystalline form. 

Besides these two forms of sulphur, prismatic crystals and 
soft sulphur, there is another. Native sulphur is crystallized 
in orthorhombic forms, generally pyramids modified by other 
forms, and the crystals obtained from solutions in carbon di- 
sulphide belong to the same system. The monoclinic needles 
in our crucible also gradually change into this form : upon 
standing they become opaque and brittle owing to a rearrange- 
ment of the molecules. 

Because sulphur has two distinct crystalline forms, it is said 
to be dimorphous. The density of prismatic sulphur is 1.96; 
that of octahedral sulphur is 2.05. At high temperatures 
the vapor density of sulphur is nearly 32, which corresponds to 
the formula S 2 . 

96. Sulphur takes fire in the air at a temperature below red- 
ness : its combustion is its union with oxygen, forming sulphur 
dioxide, SO 2 , called also sulphurous oxide and sulphurous acid 
gas. By the aid of heat, sulphur unites directly with many other 
elements : we have seen in one of our experiments (§ 4) that cop- 
per burns brilliantly in sulphur vapor, and in the same manner we 
might burn some iron wire, forming iron sulphide. 

Sulphur is used in the manufacture of matches, gunpowder, 
sulphuric acid, and many other operations. 

97. Sulphides. — We put a little antimony sulphide into a test- 
tube, and boil it with some hydrochloric acid. A gas having the 
unpleasant odor of rotten eggs is given off, and antimony chloride 
remains in the tube. This gas, which we shall now study, is called 
hydrogen sulphide, or sulphuretted hydrogen ; nearly all the sul- 
phides form this gas when boiled with hydrochloric acid, and the 
reaction gives us a test for the sulphides. 

98. Hydrogen Sulphide, H 2 S. — Into a bottle like that which 



76 



LESSONS IN CHEMISTRY. 



served for the preparation of hydrogen, we put some ferrous sul- 
phide, which we have made by heating to redness in an earthen 
crucible a mixture of iron filings with about its own weight of 
sulphur. We then pour through the funnel-tube some dilute sul- 
phuric acid, and at once or in a few minutes an effervescence shows 
us that gas is being given off, and we soon detect this gas by its 
odor. It is a compound of hydrogen and sulphur, and is formed 
by the reaction 

FeS + IPSO* - H 2 S + FeSO* 

Ferrous sulphide. Sulphuric acid. Hydrogen sulphide. Ferrous sulphate. 

The ferrous sulphate formed remains dissolved in the water. 

As we often desire this gas in the laboratory, we sometimes 
employ a self-regulating apparatus consisting of two bottles which 

__Jg__ have openings near the 
bottom, and these open- 
ings are connected by a 
stout rubber tube (Fig. 
43). In one we put a 
layer of clean pebbles 
that rise above the lower 
opening, and on this the 
ferrous sulphide; to the 
neck of this bottle we 
adapt a glass stop- cock by 
the aid of a good cork. 
In the other bottle, which 
we must not cork, we pour our dilute sulphuric acid. When we 
open the stop-cock, the acid runs in on the ferrous sulphide; the 
lias is then formed, and we may keep it in the bottle or use it as 
we desire : when we close the stop-cock, the gas forming in the 
bottle forces the acid into the other bottle, and as soon as the sur- 
face of the acid is below the top of the layer of pebbles, the fer- 
rous sulphide is no longer acted on. We may use this apparatus 
for the preparation of hydrogen and carbon dioxide (§ 234), of 
course cleaning it out before changing the materials. 

99. Properties. — Hydrogen sulphide is a colorless gas, having 




Fig 



HYDROGEN SULPHIDE. 77 

a stinking and penetrating odor. Its density being 17 times that 
of hydrogen or 17 X .0693 that of the air, its molecular weight 
must be 34 (§ 48). By strong pressure it is converted into a color- 
less liquid. At ordinary temperatures water dissolves about three 
times its volume of hydrogen sulphide, and the solution is some- 
times used in the laboratory, but it does not keep long, for the air 
oxidizes the hydrogen, forming water, while sulphur is deposited. 

We can easily determine the composition of this gas. Into a long test-tube 
of hard glass we thrust a roll of tin foil, and fill the tube with hydrogen sul- 
phide; after tightly corking the tube, we heat it until the tin acquires a yel- 
low color. After the tube has cooled, we uncork it under the surface of mer- 
cury, and we find that the volume of gas has not changed. The remaining 
gas is hydrogen, and two volumes (one molecule) of hydrogen sulphide there- 
fore contain two volumes (two atoms) of hydrogen. Subtracting, now, the 
molecular weight of hydrogen from that of hydrogen sulphide, we have 34 — 2 
= 32, the amount of sulphur in one molecule, and this we know represents 
half a molecule or one atom of that element. The formula of the gas is 
therefore H 2 S. 

Hydrogen sulphide is a combustible gas, as we can easily under- 
stand, since its molecule contains only hydrogen and sulphur, both 
of which are able to unite with the oxygen of the air, the first to 
form water, and the second to form sulphur dioxide, the same gas 
which is formed when sulphur burns in the air. When we light 
the gas at the end of the delivery-tube, it burns with a blue flame. 

100. Certain reactions of this gas make it exceedingly valuable 
in the laboratory. We pass the delivery-tube from our apparatus 
into a solution of copper sulphate in water : a brownish-black pre- 
cipitate is formed as soon as the gas comes in contact with the 
liquid. This is copper sulphide, and sulphuric acid remains in the 
solution. 

CuSO* + H2S = CuS + IPSO* 

Copper sulphate. Copper sulphide. Sulphuric acid. 

We pass the gas into a solution of antimony chloride, and an 
orange-colored precipitate of antimony sulphide forms, while hydro- 
chloric acid is in the liquid. 

2SbCl 3 + 3H2S = Sb 2 S 3 + 6HC1 

Antimony chloride. Antimony sulphide. 

In a solution of zinc acetate, we would have thrown down a 
white precipitate of zinc sulphide. Naturally, in these reactions 



78 LESSONS IN CHEMISTRY. 

we must know by analysis the composition of the molecules which 
react together, and that of the bodies which are formed, before we 
can write the equations. The solutions of many other metallic 
compounds are decomposed in this manner by hydrogen sulphide, 
and the color and other properties of the metallic sulphide formed 
show us what metal exists in the solution to which we apply the 
test. 

Hydrogen sulphide is at once decomposed by chlorine, hydro- 
chloric acid being set free. 

H2S + CI 2 = 2HC1 + S 

Hydrogen sulphide is a poisonous gas, and must not be inhaled 
for any length of time, even when very much diluted with air. 

101. Hydrosulphides. — We have seen that hydroxides are formed by the 
action of water on the oxides (J 62),- and that these hydroxides contain the 
group of atoms OH, which we call hydroxyl. On examining the composi- 
tion of the molecule of hydrogen sulphide, we see that it is analogous to that 
of water, but contains a sulphur atom instead of an oxygen atom. 
HOH HSH 

Water. Hydrogen sulphide. 

There are also compounds analogous to the hydroxides, but containing sul- 
phur instead of oxygen, and they are called hydrosulphides. We pass hydro- 
gen sulphide into a solution of potassium hydroxide; it is absorbed, and a 
chemical reaction which takes place causes the liquid to become warm. 

KOH -t HSH = KSH + HOH 

Potassium hydroxide.' Hydrogen sulphide. Potassium hydrosnlphide. Water. 

We cannot fail to notice the similarity between the structure of these mole- 
cules, and this similarity leads us to the conclusion that as far as combining 
with atoms of hydrogen and potassium is concerned, there must be a resem- 
blance between sulphur atoms and oxygen atoms. We will in time notice that 
this resemblance does not stop here, but is borne out in the structure of many 
other molecules containing sulphur and oxygen atoms. Since one atom of 
sulphur or one of oxygen is capable of combining with two atoms of hydrogen 
or one of hydrogen and one of potassium, and since we take the hydrogen 
atom as the unit of the combining power, we say that the hydrogen and 
potassium atoms are mono-atomic, and that the oxygen and sulphur atoms in 
these compounds are diatomic. 



SULPHUR DIOXIDE. 79 

LESSON XII. 
SULPHUR DIOXIDE.— SULPHUR TRIOXIDR 

102. Sulphur Dioxide, SO 2 . — This compound is formed when 
sulphur burns in the air or in oxygen : we could not obtain it 
pure by burning sulphur in air, for it would then be mixed with 
the other constituents of the air. We usually prepare the gas by 
boiling sulphuric acid with copper clippings : the products of the 
operation are cupric sulphate, water, and sulphur dioxide : know- 
ing this, we may write our equation, 

Cu + 2H2SO* = CuSO* + 2H 2 - SO 2 

Copper. Sulphuric acid. Cupric sulphate. Sulphur dioxide. 

We conduct the experiment in an apparatus like that in which 
we prepared chlorine, and if we desire to collect the gas we do so 
by downward dry displacement. 

103. Properties. — Sulphur dioxide is a colorless, suffocating 
gas. Its density compared to hydrogen is 32, agreeing with that 
which our theory should indicate (§ 48), and it is therefore a 
little more than twice as heavy as an equal volume of air. By 
pressure, or by a temperature of — 8°, it is converted into a 
colorless liquid, and this liquid may be easily prepared by passing 
the gas into a bottle surrounded by a mixture of ice and salt. 
The evaporation of the liquid produces great cold : a temperature 
as low as — 73° has been obtained by aiding this evaporation by 
pumps, and the phenomenon has been applied in the construction 
of certain machines for making ice. We can easily freeze some 
water in a test-tube by wrapping the lower end of the tube in 
some cotton wool, and pouring on this some liquid sulphurous 
oxide, but we must make the experiment in a current of air to 
carry off the suffocating gas. 

At ordinary temperatures water dissolves about forty times its 
volume of sulphurous oxide, and the solution is frequently em- 
ployed in the laboratory. 



80 LESSONS IN CHEMISTRY. 

104. Sulphurous oxide is naturally not combustible, for the 
sulphur which it contains has had an opportunity to combine 
with all of the oxygen with which it would unite. It extin- 
guishes burning bodies. 

While, however, one atom of sulphur will not combine directly 
with more than two atoms of oxygen, sulphurous oxide can be 
still further oxidized by certain reactions. If it be mixed with 
oxygen, and the mixture passed through a red-hot tube containing 
platinum sponge, the two gases combine, forming sulphur trioxide, 
SO 3 ; the vapor of this substance may be condensed by passing it 
into an ice-cold receiver. By the action of nitric acid, sulphur 
dioxide is converted into sulphuric acid, and the reaction is applied 
in the manufacture of the latter acid. 

In a tall jar we dissolve some potassium permanganate in water ; 
this body contains a large proportion of oxygen, with which it 
parts easily to oxidizable matters. We pass some sulphur dioxide 
through the purple solution, which is rapidly decolorized; the 
sulphur dioxide becomes , sulphuric acid in this reaction, and 
the potassium permanganate is said to be reduced. We use the 
term reduction to mean taking away oxygen, and any body which 
is capable of removing oxygen from other substances is called a 
reducing agent. 

Sulphur dioxide is used for bleaching wool, straw, and other 
matters which would be injured by chlorine. The substances are 
bleached by being put in a room in which sulphur is burned. We 
may in this manner bleach a flower in a bell-jar under which some 
sulphur is burning. 

105. Sulphites. — When sulphur dioxide dissolves in water, the two sub- 
stances really combine, and, by a reaction analogous to that which formed 
hypochlorous acid, sulphurous acid is formed. In this case, however, the two 
atoms of hydrogen exist in one molecule of the resulting acid. 

SO 2 + H 2 = IPSO 3 

Sulphur dioxide. Sulphurous acid. 

Sulphurous acid is not a stable compound ; it is decomposed when we try to 
separate it from its solution, and yields again sulphur dioxide and water. 
However, both of the hydrogen atoms are replaceable by metal, forming salts 
which are called sulphites ; by passing sulphur dioxide into a solution of so- 
dium hydroxide, sodium sulphite and water are formed. 



SULPHUR TRTOXIDE. 81 

2NaOH + SO 2 = Na 2 S0 3 + H20 
Sodium hydroxide. Sodium sulphite. 

To a little of this sodium sulphite in a test-tube we add hydrochloric acid; 
we can at once detect the pungent odor of sulphur dioxide, and a solution of 
common salt remains in the tube. 

Na 2 S0 3 + 2HC1 = SO 2 + H 2 -'- 2NaCl 
This gives us a test by which we may recognize a sulphite. 

106. When a sulphite is boiled with sulphur, the latter is dissolved, and a 
compound called a thiosulphate, or formerly named hyposulphite, results. 
With sodium sulphite, we would have sodium thiosulphate. 

Na 2 S0 3 + S = Xa 2 S 2 3 or Na 2 S0 3 S 

Sodium sulphite. Sodium thiosulphate. 

It will be noticed that the thiosulphate has exactly the composition of a 
sulphate (§ 113) in which an atom of oxygen is replaced by an atom of sul- 
phur. When a thiosulphate is treated with an acid, sulphur dioxide is 
evolved, and sulphur separates. The rags used in the manufacture of paper 
are bleached by chlorine, but no chlorine must be left in the paper, or this 
would be injured. In presence of water, sulphur dioxide (we may then say 
sulphurous acid) and chlorine react to form sulphuric and hydrochloric acids, 
both of which may readily be neutralized. 

H 2 S0 3 + CI 2 + H 2 = H 2 30* + 2HC1 
Sulphurous acid. Sulphuric acid. 

Sodium thiosulphate therefore serves as an antichlor in the manufacture or 
paper. 

107. Sulphur Trioxide, SO 3 . — We have seen that this com- 
pound is formed by the direct union of sulphur dioxide and oxy- 
gen in presence of heated pla:inum sponge (§ 104) : other porous 
substances cause the same combination. Sulphur trioxide is usu- 
ally prepared by heating fuming sulphuric acid, which is some- 
times called Nb-rdhausen acid, because it was for a time manufac- 
tured only in the village of Nordhausen, in the Hartz, by distilling 
partially-dried ferrous sulphate. It contains a compound of 
sulphur trioxide and sulphuric acid ; H 2 SO* + SO 3 = H 2 S ? 7 . 
When this is heated, it decomposes into its constituents ; the 
sulphur trioxide, being the most volatile, is condensed in cold 
flasks, which are at once hermetically sealed. 

Sulphur trioxide is a snowy-white solid, crystallizing in feather- 
like flakes. It combines so energetically with water that each 
particle makes a hissing noise like hot iron on touching the liquid. 
The result of this combination is sulphuric acid. 
SO 3 + H 2 = IPSO 



82 LESSONS IN CHEMISTRY. 

LESSON XI1L 

SULPHURIC ACID, H 2 SO*. 

108. Into a jar of oxygen containing a little water, we lower a 
deflagrating-spoon containing burning sulphur. The jar soon be- 
comes filled with sulphur dioxide, and when the flame of the sul- 
phur is extinguished, we pour into the jar a little nitric acid. Red 
vapors at once become apparent, but disappear in a little while: 
in order to mix the gases well, we now shake the jar, keeping it 
closely covered, and then by means of a long glass tube we blow 
some air into it : red vapors are again produced. It is evident 
that some chemical change has occurred between the nitric acid 
and sulphur dioxide, and that another change takes place between 
the gases in the jar and the air which we have introduced. The 
first change is the production of sulphuric acid, and the conversion 
of the nitric acid into the red vapors of nitrogen peroxide. Since 
one molecule of nitric acid contains only one atom of hydrogen, 
while one molecule of sulphuric acid contains two such atoms, two 
molecules of nitric acid must react with one of sulphur dioxide. 

SO 2 + 2HN0 3 = H 2 SO* + 2N0 2 

Nitric acid. Sulphuric acid. Bed vapors. 

The red vapors react with the water in the jar and more of the 

sulphur dioxide, converting the latter into sulphuric acid. 

SO 2 + NO 2 + H 2 = H 2 SO* 4- NO 

Nitric oxide. 

But the nitric oxide, which is a colorless gas, is not lost : it 
takes an atom of oxygen from the air blown into the jar, and 
again forms red vapors. 

NO + = NO 2 
These red vapors in turn react with water and more sulphur 
dioxide, and this series of reactions continues until all of the 
sulphur dioxide is converted into sulphuric acid. Thus, a small 
quantity of the oxides of nitrogen will effect the conversion of 



SULPHURIC ACID. 83 

a large quantity of sulphur dioxide into sulphuric acid, the 
oxygen needed being taken from the air. 

109. These reactions are those which actually take place in the 
manufacture of sulphuric acid, which is commonly called oil of 
vitriol, and of which enormous quantities are used at one stage or 
another in the manufacture of nearly all other chemical com- 
pounds. The sulphur dioxide is obtained by burning sulphur in 
furnaces (A A, Fig. 44), the heat of which boils water for the 
steam required in the operation. The sulphur dioxide formed 
passes through a series of leaden chambers, in one of which (D) 
it comes in contact with nitric acid, which trickles down over a 
sort of cascade (EE). The gases then pass through other leaden 
chambers into which steam is injected (HH) : a further amount 
of sulphuric acid is produced, and this collects on the floor 
of the chambers, from which it is drawn off. An excess of air 
must be passed into the chambers in order to reoxidize the 
nitric oxide, and as the nitrogen of the air which must be allowed 
to escape from the apparatus would carry off some of that oxide 
of nitrogen, all the waste gases are obliged to pass through a 
tower (R) filled with coke which is kept wet with strong sulphuric 
acid. This latter absorbs the nitrous gases, and as it runs from 
the tower is conducted into a vessel (i), from which it may be 
forced by steam pressure to the top of the first small chamber (C), 
through which the sulphur dioxide is caused to pass. Here the 
sulphur dioxide removes all of the nitrous gases and carries them 
again into the chambers, so that only nitrogen from the air used 
escapes at the chimney of the coke column. The chambers are 
of various sizes, sometimes five metres wide and high, and ten, 
twenty, or even more metres in length. 

The acid drawn from the leaden chambers is called chamber 
acid: it is strong enough for many purposes, its density being 
1.5. The strong acid, density 1.842, is made by evaporating 
the chamber acid in leaden boilers until its further concentration 
would dissolve the lead ; it is then transferred to expensive plati- 
num stills, where the evaporation is terminated. 

110. The sulphuric acid of commerce always contains a little 



84 



LESSONS IN CHEMISTRY. 




SULPHURIC ACID. 85 

lead sulphate, formed in the chambers and evaporating boilers, 
and when it is diluted with water this lead sulphate becomes in- 
soluble and separates as a white precipitate. It is often brown 
from the presence of a little carbonaceous matter. Sometimes the 
sulphur dioxide is obtained by burning iron pyrites (iron disul- 
phide), and, as the pyrites often contaios arsenic, the resulting sul- 
phuric acid also contains arsenic. Pure sulphuric acid is made 
by distilling the commercial acid in glass retorts ; the operation 
requires great care, for the retorts sometimes break, and the vapors 
of the sulphuric acid are most corrosive and suffocating. 

111. Properties. — Pure sulphuric acid is a colorless, oily liquid, 
having at 12° a density of 1.842. It solidifies at 10.5°, and boils 
at about 338° : its boiling is accompanied by explosive emission of 
vapor, which may be obviated by putting some pieces of platinum 
in the vessel. Sulphuric acid is soluble in all proportions of water, 
and the mixture is accompanied by the production of great heat, 
showing that there is a true chemical combination between the 
water and acid. In diluting sulphuric acid with water, we always 
pour the acid very gradually into the water, which we stir con- 
stantly. If the mixture is made suddenly, part of the acid is 
sometimes thrown out of the vessel. 

The affinity of sulphuric acid for water is so strong, that the 
acid causes the formation of water in many substances which do 
not contain water, but contain hydrogen and oxygen in the pro- 
portions required for its formation. In a beaker glass, or other 
thin glass vessel, we pour a little strong solution of sugar, and 
then some concentrated sulphuric acid : instantly the mixture turns 
black, and a mass of porous charcoal fills the vessel, which may 
overflow if we have used too much of the materials. Sugar con- 
tains carbon or charcoal, and oxygen and hydrogen in the propor- 
tions to form water. For the same reason a chip of wood with 
which we stir some sulphuric acid quickly becomes blackened, and 
the brown color which the acid acquires shows us why the common 
acid is often brown. 

When sulphuric acid is passed through a red-hot tube, it is 
decomposed into sulphur dioxide, oxygen, and water. 



86 LESSONS IN CHEMISTRY. 

H 2 SO* = SO 2 + + H20 
We have already seen how zinc acts on sulphuric acid, replacing 
the hydrogen which then becomes free. The action of copper on 
the acid is a reducing action, part of the sulphuric acid being 
reduced to sulphur dioxide. 

112. Molecular structure of sulphuric acid. — We have seen that hypochlorous 
acid contains the group hydroxyl, OH; sulphuric acid also contains this group, 
and we may understand the structure of its molecule by studying some simple 
reactions. Since a molecule of sulphur dioxide contains two atoms of oxygen, 
each of which is diatomic, — that is, capable of combining with two atoms of 
hydrogen, — the sulphur atom must in this compound have as much combining 
power as four atoms of hydrogen : we call it tetratomic. Yet this sulphur atom 
is capable of combining with another atom of oxygen ; it is unsaturated with 
oxygen, although it is satisfied with the two atoms. We mix in a glass jar 
equal volumes of chlorine and sulphur dioxide, and expose the mixture to 
direct sunlight : the gases combine to form a colorless liquid, having a suffo- 
cating vapor, and we call the compound svlphuryl chloride. Analysis shows 
that it contains S0 2 C1 2 , and for convenience' sake the group of atoms SO 2 is 
called sulphuryl. Each atom of chlorine is worth one of hydrogen, and if 
in sulphur dioxide the sulphur atom is tetratomic, it must be hexatomic in 
sulphuryl chloride. We may represent this relative combining capacity by 
little lines, which will show us, not how and where the atoms unite, but the 
relative worth of the atoms in combination. The free atom of hydrogen or 
of chlorine would be indicated to be monatomic by a single line, thus, H-, C1-; 
and in the molecules of these two elements or of their one compound a single 
line between the two symbols would show that one has as much combining 
power as the other : 

H-H Cl-Cl H-Cl 

The symbol of a diatomic element must have two lines, to show that it is 
worth two monatomic atoms, and we may write water and hydrogen sulphide, 

H-O-H H-S-H 

For sulphur dioxide, sulphur trioxide, and sulphuryl chloride, in the first 
of which the sulphur atom is tetratomic and in the other two hexatomic, we 
must show the combining power, or atomicity, as it is often called, of the 
elements by four or six lines, and show that the oxygen is diatomic by giving 
its atoms each two lines. We therefore write, 

0=S=0 CI 

i w // 

0=S=0 8 

ci 6 

Sulphur dioxide. Sulphuryl chloride. Sulphur trioxide. 

When sulphuryl chloride is poured into water, both substances are decom- 
posed, sulphuric and hydrochloric acids being formed. 

S0 2 C1 2 + 2H 2 = H 2 S0 4 + 2HC1 



SULPHATES. 87 

We must explain this reaction by a replacement of the chlorine atoms in the 
sulphuryl chloride by other atoms, or groups of atoms, which have the same 
combining power, and we would then conclude that sulphuric acid contains 
two hydroxyl groups, and we might call it sulphuryl hydrate. 



Cl-S-Cl + H-O-H + H-O-H = H-O-S-O-H + HC1 + HC1 


Sulphuryl Water (two molecules). Sulphuric acid. Hydrochloric acid 

chloride. (2 molecules.) 

Analogous study of the manner in which compounds are formed and decom- 
posed, and of the relative worth of the atoms, has led chemists to hold definite 
ideas of the relations which the atoms bear to each other in a great number 
of molecules. The group of atoms OH is called a compound radical because it 
is capable of replacing an atom or simple radical. In the same manner the 
group SO 2 is a radical, as in general are all groups of atoms which pass by 
double decomposition and without change from one molecule to another. Some 
radicals, like SO 2 , can be separated and studied, because the combining power 
of the atoms in them, though not saturated, is satisfied ; others, like -OH, can- 
not be separated, because their atoms are not satisfied with each other. Why 
this is, chemists have not yet been able to explain satisfactorily, but the fact 
may be stated that a monatomic radical cannot usually exist except in com- 
bination : this applies to monatomic atoms as well as to monatomic compound 
radicals ; we have already seen that the molecules of hydrogen and chlorine 
must each contain two atoms. 

There are two other elements whose atoms exactly resemble sulphur in their 
power of combining with other atoms. They are selenium and tellurium. 
They are found only in small quantities, usually associated with gold and 
silver ores. 



LESSON XIV. 

SULPHATES. 



113. When the hydrogen of sulphuric acid is replaced by 
metals, sulphates are formed, but as there are two atoms of hydro- 
gen, and either one or both may be replaced, we can understand 
that there may be two kinds of sulphates. If only one hydrogen 
atom be replaced, the resulting salt will have acid properties, for 
it still contains an atom of replaceable hydrogen ; but if both be 
replaced, we have a neutral salt, — that is, one which is neither acid 



88 LESSONS IN CHEMISTRY. 

nor alkaline. We may study the formation of two of these salts 
in the reaction of one and two molecules of sodium hydroxide with 
one molecule of sulphuric acid. 

H2SO + NaOH = NaHSO* + IPO 

Sodium hydroxide. Sodium acid sulphate. 

H 2 S0 4 + 2NaOH = Na2SO* + 2H20 
Sodium sulphate. 

But in the action of zinc on sulphuric acid, one atom of zinc 
replaces two atoms of hydrogen. In the same manner, if we boil 
sulphuric acid with lead oxide*, we have formed lead sulphate, in 
which one atom of lead replaces both atoms of hydrogen. 

PbO + H' 2 SO* = PbSO* + H20 

Lead oxide. Lead sulphate. 

Since one atom of zinc or one of lead is thus capable of replacing 
two atoms of hydrogen, those metals are said to be diatomic ; and 
since sulphuric acid contains two atoms of hydrogen which may be 
replaced by two atoms of a monatomic metal, like sodium, or by one 
atom of a diatomic metal, like zinc, it is called a dibasic acid, and 
is capable of forming neutral and acid salts. 

114. With the exception of the sulphates of barium, strontium, 
and lead, all the sulphates are soluble in water, but calcium 
sulphate, silver sulphate, and mercurous sulphate are only 
slightly soluble. 

115. To a solution of magnesium sulphate we add a few drops 
of solution of barium chloride or barium nitrate. A white cloud 
forms ; this is insoluble barium sulphate ; when it has settled, we 
may pour off most of the liquid, and we will find that our white 
substance is not dissolved by boiling nitric acid. This test enables 
us to recognize either a soluble sulphate or uncombined sulphuric 
acid. 

Some of the sulphates form anhydrous crystals, — that is, with- 
out water; others require water of crystallization. 

116. Sodium Sulphate, Na 2 S0 4 , was for a long time called 
Glauber's salt, because Glauber found that it was useful as a 
purgative medicine. It crystallizes in colorless, oblique rhombic 
prisms containing ten molecules of water of crystallization, so that 



SULPHATES. 89 

the formula of the crystals is Na 2 S0 4 + 10H 2 O. They are sol- 
uble in about ten times their weight of water at 0°, and in one- 
third their weight at 33° ; if a saturated solution be made at the 
latter temperature and immediately sealed, it will remain liquid 
indefinitely, but on opening the flask the whole of the liquid in- 
stantly becomes a mass of crystals. 

117. Potassium Sulphate, K 2 S0 4 . forms very hard, colorless 
crystals, not very soluble in water ; it is poisonous. 

118. Calcium Sulphate, CaSO 4 . — We have in a beaker glass a 
very strong solution of calcium chloride. To this we add at arm's 
length about half its volume of concentrated sulphuric acid : the 
contents of the beaker at once become so solid that we can in- 
vert it and nothing runs out. This solid is calcium sulphate. 

CaCl 2 + H 2 SO± = CaSO* 4- 2HC1 
Calcium chloride. Calcium sulphate. 

Calcium sulphate is the mineral gypsum, alabaster, or selenite. 
In these minerals it is combined with two molecules of water of 
crystallization ; this water is driven out when they are heated to 
120°, leaving the anhydrous sulphate as a fine white powder, 
known as plaster of Paris. Unless it has been heated to too high 
a temperature, this substance will again combine with its water of 
crystallization, and such combination takes place when plaster of 
Paris is mixed with water. Plaster casts are made by mixing the 
plaster and water to a creamy consistence, and pouring the liquid 
into the moulds : in a few minutes the plaster sets, or becomes 
hardened, 2nd in so doing it expands and completely fills the 
mould. Calcium sulphate dissolves in about 500 times its weight 
of water. It is a valuable fertilizer for certain soils. 

119. Strontium Sulphate, SrSO 4 , constitutes the mineral celes- 
tine, so called because it often has a blue color, though the pure salt 
is white. It is insoluble in water, and is precipitated when a soluble 
strontium salt is added to sulphuric acid or a soluble sulphate. 

120. Barium Sulphate, BaSO, is found native as heavy spar. 
We have seen that it is formed by the reaction of sulphuric acid 
with soluble salts of barium. It is sometimes used for adulterating 
white lead (§ 250). 



90 LESSONS IN CHEMISTRY. 

121. Magnesium Sulphate, MgSQ 4 + 7H 2 0, is commonly 
known as Epsom salts. It is made by dissolving magnesium car- 
bonate in dilute sulphuric acid, and when the concentrated solu- 
tion is allowed to evaporate, the salt separates in crystals containing 
seven molecules of water. It has a salty, bitter, and unpleasant 
taste. It dissolves in about three times its weight of water. It is 
used in medicine. 

122. Zinc Sulphate, ZnSO 4 + 7H 2 0.— We evaporate to a 
small volume the liquid remaining in the bottle in which we made 
hydrogen by the action of sulphuric acid on zinc, and then set it 
aside in a cool place. After a time zinc sulphate separates in 
beautiful transparent crystals, containing seven molecules of water, 
and of exactly the same form as those of magnesium sulphate 
prepared in the same manner. Compounds which have in their 
molecules the same number of atoms arranged in the same man- 
ner, usually crystallize in the same form, and are said to be iso- 
morphous. Zinc sulphate is sometimes called white vitriol. It 
is quite soluble in water, and when swallowed it acts as a violent 
emetic. 

123. Ferrous Sulphate, FeSO 4 + 7H 2 0.— This salt, called 
also green vitriol and copperas, is made by treating scrap iron 
with dilute sulphuric acid ; hydrogen is disengaged, just as in the 
action of the same acid on zinc. When the filtered solution is 
evaporated and set aside to crystallize, the ferrous sulphate 
separates in monoclinic crystals which contain seven molecules 
of water of crystallization. These crystals are pale green in 
color ; when exposed to dry air, they lose part of their water 
of crystallization, and the surface becomes covered with a white 
powder, which is the anhydrous salt ; they are said to effloresce. 
After a time this powder becomes yellow, from an absorption of 
oxygen (§ 527). Ferrous sulphate is soluble in less than twice its 
weight of cold water, and much more soluble in boiling water. 
It is poisonous. 

124. Cupric Sulphate, CuSO 4 + 5H 2 0.— This beautiful blue 
salt, often called blue vitriol, may be prepared from the res- 
idue of the preparation of sulphur dioxide by diluting it with 



NITROGEN. — THE ATMOSPHERE. 91 

water, filtering, and evaporating to crystallization. It is usually 
made by roasting — that is, heating in the air — copper sulphide 
(§ 484), and treating the mass with water. When the blue crys- 
tals are heated, the water is driven out, and the white anhydrous 
salt is left. Cupric sulphate dissolves in four times its weight of 
cold, or twice its weight of boiling water. To a solution of this 
salt we add a little ammonia water ; a pale-blue precipitate forms, 
but when we add more ammonia this precipitate again dissolves, 
and a deep-blue liquid is obtained. This liquid contains ammo- 
itiacal cupric sulphate. Cupric sulphate is used in telegraphic 
batteries, in dyeing, for electrotyping, and in many other opera- 
tions. It is poisonous. 

125. Lead Sulphate, PbSO 4 . — When sulphuric acid or the 
solution of a sulphate is added to a solution containing a lead salt, 
lead sulphate separates as a white precipitate. It occurs in nature 
as the mineral anglesite. It is insoluble in water, but dissolves 
in strong acids. 

Of the many other sulphates, we must study a few when we 
shall have learned some of the peculiarities of the corresponding 
metals. 



LESSON XV. 
NITROGEN.— THE ATMOSPHERE. 

126. Nitrogen, N = 14. — On the water in the pneumatic 
trough, we float a small capsule containing a little sand on which 
we have placed a piece of phosphorus. We ignite the phos- 
phorus, and place over it a bell-jar which may rest on the shelf 
in the trough (Fig. 45). At first, as the heat of the burning 
phosphorus expands the air, a few bubbles of air escape under the 
edge of the jar, but this soon stops; presently the water begins 
to rise in the jar, and the phosphorus no longer burns. All 
the oxygen of the air in the jar has been consumed by the phos- 
phorus, and there is left nitrogen, with which the oxygen was 



92 



LESSONS IN CHEMISTRY. 




FlQ. 45. 



mixed, and phosphoric oxide. The latter will presently dis- 
solve in the water, and we may then examine the nitrogen, 
which contains but very small quantities of other gases. 

To prepare larger quantities of nitrogen free from most of 
these impurities, we pass a current of 
air through a tube containing pieces of 
solid potassium hydroxide, which absorbs 
the moisture and carbon dioxide, and 
then through a long tube containing red- 
hot copper. The copper combines with 
jgljb, the oxygen, forming cupric oxide, and 
g^ nitrogen passes out at the end of the 
tube. Perfectly pure nitrogen can be 
obtained only by decomposing certain 
compounds, such as ammonium nitrite. 
Upon heating, this salt yields nitrogen 
and water, NH 4 N0 2 = N 2 + 2H 2 0. 

127. Nitrogen is a colorless, tasteless, and odorless gas. Its 
density compared to air is 0.97, or compared to hydrogen, 14 ; 
as its atomic weight is also 14, its molecule must contain two 
atoms. It is almost insoluble in water, and difficult to liquefy. 
It is not combustible, neither will it support the combustion of 
other substances. It combines directly with only a few of the 
elements, and energy is absorbed during the formation of many 
of its compounds ; that is, the nitrogen atoms have a stronger 
affinity for one another than for the other atoms with which 
they are combined. 

128. The Atmosphere. — The chemical composition of the air 
was first determined with tolerable accuracy by the great French 
chemist, Lavoisier. We may satisfy ourselves of this composition 
in a very simple manner. Around a long glass tube, closed at 
one end, we have placed four caoutchouc bands, dividing it into 
five equal portions. Into this tube, which must be perfectly dry, 
we drop a dry piece of phosphorus, and tightly cork the open 
end. By gently heating the bottom of the tube over a lamp, we 
inflame the phosphorus, and then by quickly turning the tube 



THE ATxMOSPHERE. 



93 



bottom up and giving with the corked end a few sharp blows on 
the table, we cause the burning phosphorus to fall the whole 
length of the tube. If our experiment has been well made, all 
the oxygen has been burned from the air in the tube, which we 
allow to cool, and then carefully uncork with 
the mouth under water. As soon as the cork 
is drawn, the water rises to the first division 
(Fig. 46). The air which we have roughly 
analyzed, then, contained about one- fifth oxy- 
gen and four-firths nitrogen by volume. 

129. A very accurate analysis of air is 
made by the aid of the eudiometer, which we 
have studied. Into the eudiometer with the 
caoutchouc tube and plain glass tube, which 
served for the synthesis of water (§ 42), but 
without the enclosing wide glass tube, we in- 
troduce 100 measures of air and 100 meas- 
ures of pure hydrogen. After adjusting the 
mercury level in the two tubes, we pass an 
electric spark : at once the oxygen and part 
of the hydrogen are converted into water, 
which condenses, and the volume of gas is re- 
duced. We know that water is formed by the union of two 
volumes of hydrogen and one volume of oxygen ; consequently 
one- third of the diminution in volume must be caused by the 
removal of the oxygen of the 100 measures of air. On again 
adjusting the level of the mercury, we find that instead of 200 
measures we have only 137.21. The oxygen present in 100 

200 — 137.21 




Fig. 46. 



measures of air must, then, have been 



or 20.93 



measures. We conclude, therefore, that 100 volumes of air con- 
tain 20.93 volumes of oxygen and 79.07 volumes of nitrogen.* 
Since oxygen is heavier than nitrogen, these relative volumes 
will not express the relations by weight. We can calculate the 
weights from the volumes, and the result would show us that 76.87 
parts by weight of nitrogen are mixed with 23.13 parts of oxygen. 



* This nitrogen contains a small amount of argon (jf 130). 



94 LESSONS IN CHEMISTRY. 

These proportions are confirmed by the direct analysis, which is made by 
passing air through a series of tubes in which all traces of carbon dioxide 
and moisture are absorbed : thus purified, the air passes through a tube con- 
taining red-hot copper ($ 127), and the increase in weight of this tube gives 
the amount of oxygen in the air analyzed. The nitrogen passes on into a 
glass globe in which a vacuum has previously been made, and of course the 
increased weight of this globe is the amount of nitrogen. 

The air is not a compound, but a mixture, and we may expect 
that the proportions of the constituents shall vary a little. How- 
ever, the composition is nearly constant : hundreds of analyses 
have shown that the proportion of oxygen in 100 volumes of 
unconfined air varies only from 20.86 to 21. 

130. Argon.— The nitrogen remaining after the absorption 
of oxygen from purified air is not pure. It was observed quite 
recently that it is slightly heavier than the pure nitrogen ex- 
tracted from chemical compounds, and further shown that this 
difference in density is due to a hitherto unknown element, 
which was called argon. When " atmospheric" nitrogen is 
passed over heated magnesium, the nitrogen gradually combines 
with the metal, and a small volume of unabsorbable gas re- 
mains. This has a density of 19.9. It is rather more soluble 
in water than nitrogen, and has been reduced to both the liquid 
and the solid states at very low temperatures. It is most re- 
markable for its entire lack of chemical affinities : all attempts 
to combine it with other elements have resulted negatively, 
hence the name argon (apyov, inactive). We have reason to 
believe that each molecule consists of a single atom. The 
atomic weight, therefore, is nearly 40. 

131. We pour into a plate some clear lime-water; in a few 
minutes a thin, white pellicle forms over its surface. This is cal- 
cium carbonate, and has been formed by the absorption of carbon 
dioxide from the air. Air also contains more or less vapor of 
water, which is deposited in the form of dew on very cold objects. 

The proportions of vapor of water and carbon dioxide may be determined 
by drawing a known volume of air through a series of tubes (Fig. 47), the 
first of which contain pumice-stone and sulphuric acid, and the others frag- 
ments of potassium hydroxide. The increase in weight of the first tubes 



THE ATMOSPHERE. 



95 



(D, E, F) gives the weight of that vapor, and the increase in weight of the 
tubes containing potash (A, B, C) gives us the proportion of carbon dioxide. 
The volume of air which contained these quantities is equal to the volume of 
water which runs from the aspirator (V). We can calculate the weight of this 
air from its volume, for at 0° and under 760 millimetres barometric pressure, 
one litre of dry air weighs 1.2932 grammes. 

A gas expands or contracts 0.00366 of its volume at 0° for every degree 
above or below 0°, and since the volume of a gas is inversely as the pressure 
(Mariotte's law), the volume of any gas may be calculated for 0° and 760 mil- 
limetres by the equation V = 75 U + o .u036 6t) WherG Fre P resents the volume 
at t°, and h the barometric pressure expressed in millimetres. 

132. The quantity of vapor of water which the air can take 
up depends on the temperature, and air is said to be saturated 
with moisture when at the given temperature it can hold no more 
water vapor. It is then said 

to have a relative humidity 
of 100 : at the same tem- 
perature half that quantity 
of vapor would be a relative 
humidity of 50. But if the 
temperature be increased 
and the quantity of moist- 
ure remain the same, the 
relative humidity is lowered, 
for the air is then capable 

of dissolving more vapor. The temperature at which air is 
completely saturated with vapor is called the dew-point, and 
this may be determined by noting the temperature at which 
moisture begins to deposit on the walls of a vessel which is 
artificially cooled. Substances which are capable of absorbing 
moisture from the atmosphere are said to be hygroscopic. Sul- 
phuric acid, for instance, when exposed to the air will in a few 
days take up enough moisture to double its volume. 

133. Carbonic acid gas is present in the air in only small pro- 
portions ; from four to six parts in ten thousand parts of air. It 
is thrown into the atmosphere from volcanoes, fissures in the earth, 
and mineral springs, but the largest quantity is produced by com- 
bustion and respiration. It does not accumulate in the atmosphere, 




96 



LESSONS IN CHEMISTRY. 




Fig. 48. 



but is absorbed by plants, and under the influence of sunlight is 
decomposed, the carbon being retained for the growth of the plant, 
while oxygen is eliminated. If we put some tender 
leaves, water-cress answers very well, in a jar which 
we fill with water charged with carbonic acid, 
and place on a plate so that the water may not run 
out, and then expose to direct sunlight, in a short 
time bubbles of gas collect in the jar (Fig. 48). 
1 We may transfer this gas to a small tube, and if 
we test it by a lighted match, we find that it is 
oxygen. We can prove that carbon dioxide ex- 
ists in the air exhaled from the lungs, by blowing the breath through 
lime-water (Fig. 49), which quickly becomes clouded by the for- 
mation of calcium carbonate. In the same manner, if we burn a 

lighted taper or candle in a 
covered jar, and then pour in 
some lime-water, and shake 
the jar, the milkiness of the 
water shows that carbon di- 
oxide has been formed. 

134. Although in uncon- 
fined air, plants and vegetables 
remove the carbon dioxide, so 
that its proportion does not 
increase, yet if the air be con- 
fined, as in a room or a mine, 
this gas may accumulate to as 
much as one part in a hundred 
of air. As this carbon dioxide 
is formed at the expense of 
the oxygen of the air, the proportion of oxygen may descend as 
low as 22 parts per hundred by weight, instead of 23.2. At 
every breath a man consumes about 4.87 per cent, of the oxygen 
which he inhales, and the carbon dioxide exhaled in an hour is 
about 20 litres. When the carbon dioxide in the air is pure, 
its proportion may be much increased, and no ill effects result ; 




Fig. 49. 



AMMONIA. 



97 



but in addition to this gas a considerable proportion of animal 
matters passes from the lungs, and, together with that thrown 
off in the perspiration, quickly vitiates the atmosphere of an 
apartment which is not ventilated. Good ventilation requires 
from six to ten thousand litres of air per hour for each indi- 
vidual. In dwellings and workshops, most of the ventilation is 
by the cracks of doors and windows. Fires in open grates afford 
excellent ventilation, the draught of the chimney drawing a 
constant supply of air into the room. 

135. Besides the substances already considered, air always con- 
tains very small quantities of ammonia, traces of nitric acid, and 
small solid particles of various natures which are carried to great 
distances by the winds. Sometimes a little ozone is present, and 
may be recognized by the test which we have studied (§ 66). 



LESSON XVI. 



AMMONIA AND ITS COMPOUNDS. 



-In a glass flask to which we have 



136. Ammonia, NH 3 
adapted a cork and 
delivery-tube, we mix 
some powdered ammo- 
nium chloride with its 
own weight of slaked 
lime. We then fill 
the rest of the flask 
with pieces of quick- 
lime, and gently heat 
it on a sand-bath. We 
soon notice the pungent 
odor of the gas disen- 
gaged ; as this gas is very soluble in water, we cannot collect it 

7 




Fig. 50. 



98 



LESSONS IN CHEMISTRY. 



over that liquid ; we may collect it either over mercury in a 
small pneumatic trough, or by upward dry displacement, for it is 
lighter than air (Fig. 50). Slaked lime is calcium hydroxide, 
Ca(OH) 2 , ammonium chloride is a compound of nitrogen, hydro- 
gen, and chlorine, NH 4 C1. We may write the reaction, 

2NH*C1 + Ca(OH)2 = 2 NH3 + CaCl 2 + 2H*0 
Ammonium chloride. Lime. Ammonia. Calcium chloride. 

The calcium chloride formed remains in the flask, and the 
water is absorbed by the pieces of lime which we have put into 
the flask for that purpose. We could not dry ammonia gas by 
passing it over either calcium chloride or sulphuric acid, for it 
combines with both of those substances. 

137. Properties. — The ammonia which we have collected is a 
colorless gas, having a penetrating, pungent odor, and a burning 
taste. We must not inhale too much of it, for, although not 
poisonous, it often produces sudden giddiness or vertigo. Its 
density compared to hydrogen corresponds with half its molecular 

weight, being 8.50 : it is therefore a 
little more than half as heavy as air. 
By strong pressure, it is readily con- 
verted into a liquid, and this liquid 
is employed in some forms of ice- 
machines, where it produces great cold 
by its evaporation. 

Ammonia is very soluble in water : 
at 0° water will dissolve 1000 times 
its volume of the gas, and at ordinary 
temperatures about 700 times its vol- 
ume. We have fitted to a glass flask 
a cork through which passes a tube 
drawn out to a small opening on the 
inside. We fill this flask with am- 
monia, by dry displacement, and after 
putting in the cork we dip the end of 
the tube into a vessel of water. The water slowly rises in the 
tube, but as soon as it reaches the narrow end the ammonia is 




AMMONIA AND ITS COMPOUNDS. 



99 



absorbed so rapidly that the pressure of the atmosphere forces the 
water up in a fountain which continues until all of the ammonia 
is dissolved (Fig. 51). The solution of ammonia in water is 
called ammonia-water, liquor ammonise, or spirits of hartshorn. 
It has the taste and odor of the gas, and is very caustic. When it 
is heated, the gas is driven out, and we may most readily obtain 
ammonia by heating strong ammonia-water in a flask, and drying 
the gas by passing it through a tube containing quick-lime. The 
strong ammonia-water of commerce contains about 35 per cent, of 
the gas. Its density is about 0.86. 

138. Ammonia is decomposed into nitrogen and hydrogen by very high 
temperatures or by the continued passage of electric sparks. Two volumes of 
ammonia yield four volumes of the mixed gases, and if we mix in the eudiom- 
eter these four volumes with one and a half volumes of oxygen and pass the 
spark, after the condensation of the water formed, only one volume of gas is 
left. This is nitrogen, and two volumes, or one molecule, of ammonia must 
therefore contain one volume (one atom) of nitrogen, and three volumes (three 
atoms) of hydrogen. 

139. Ammonia is combustible, but it will not burn in the air. 
We may cause it to burn at a jet which is surrounded by oxygen, 
and for that purpose we have fitted to a 
short wide tube, open at both ends, a cork 
through which pass two tubes (Fig. 52) ; 
one of them is short and leads oxygen 
from a gas-holder, while the other reaches 
nearly to the top of the wide tube, and 
conveys ammonia gas from a small flask 
in which we boil some ammonia-water. 
As soon as ammonia-gas escapes from 
the jet, we turn on the oxygen, and light 
the ammonia, which burns with a yellow 
flame, forming water and nitrogen. 
4NH3 + 30 2 = 6H 2 + 2N 2 

This combustion may be made to take 
place more slowly, and in an interesting manner, in the presence 
of platinum. Over some ammonia-water contained in a beaker 
glass (Fig. 53) we suspend a coil of red-hot platinum wire, so 




Fig. 52. 



100 



LESSONS IN CHEMISTRY. 




Fig. 53. 



that it may Dearly touch the liquid. The coil will continue to 
glow for a long time by the heat evolved from the slow combustion 
of the ammonia which escapes from the liquid 
and mixes with the oxygen of the air. If now 
we warm the beaker, and pass bubbles of oxygen 
through the liquid, each bubble causes a little 
explosion as it combines with the hydrogen of 
the ammonia. Sometimes the beaker becomes 
filled with white fumes of ammonium nitrite. 
140. Ammonium Compounds. — Ammonia is 
not the only compound of nitrogen and hydrogen. Hydrazine 
and hydrazoic acid are the names of two remarkable substances 
recently discovered, and which have the compositions N 2 H* and 
N 3 H respectively. 

There is also a class of compounds which contain more than 
three atoms of hydrogen for each atom of nitrogen. Over a 
small capsule containing some warm ammonia-water we have 
inverted a glass jar, and, at a little distance, over another cap- 
sule in which is some 
warm hydrochloric 
acid we have inverted 
another jar. Each 
jar now contains some 
of the gas from the 
liquid under it. When 
we raise the jars and 
bring their mouths to- 
gether, both become filled with dense white fumes (Fig. 54). 
The two gases have combined, and a body called ammonium chlo- 
ride has been formed, and will settle on the sides of the jars. The 
combination is very simply expressed. 

NH 3 + HCl = NH*C1, ammonium chloride. 
We see, then, that while the nitrogen atom will combine with 
only three atoms of hydrogen alone, it will combine with four if 
an atom of chlorine come with that hydrogen. In the same man- 
ner, in many other compounds one nitrogen atom is combined with 




AMMONIUM CHLORIDE. 



101 



four hydrogen atoms, and one other atom or group of atoms. NH 4 
is one of those groups of atoms which we call radicals (§ 112) ; it 
passes from one compound to another without change, just as an 
atom of hydrogen may pass from one molecule to another. It 
cannot, however, be separated in the free state from any of these 
compounds. It is called ammonium. 



LESSON XVII. 



AMMONIUM COMPOUNDS. 

141. Ammonium Chloride, NH 4 C1. — This compound is formed 
by the direct union of ammonia and hydrochloric acid. During the 
manufacture of illuminating gas by the distillation of coal, more 
or less ammonia is formed ; it must be removed before the gas is 
fit for use, and this is accomplished by washing the gas with water 
(§ 225). A dilute solution of ammonia is thus obtained, and this 
is the source of the ammonia and ammo- 
nium compounds of commerce. For the 
preparation of ammonium chloride this 
gas liquor is heated with lime, and the 
ammonia gas given off is passed into 
hydrochloric acid. The solution is then 
evaporated, and the residue of ammonium 
chloride is purified by sublimation in 
stoneware pots. It may be formed by 
another and interesting reaction : we pass 
into a jar of dry chlorine the drawn-out 
end of a tube through which ammonia is 
escaping ; at once the ammonia takes fire, 

being partially decomposed with production of hydrochloric acid, 

which at once unites with another portion of the ammonia, forming 

white clouds of ammonium chloride (Fig. 55). 

2NH 3 -f 3C1 2 = N* + 6HC1 

6HC1 + 6NH 3 = 6XH*.C1 




Fig. 



102 LESSONS IN CHEMISTRY. 

When pure, ammonium chloride is in translucent masses, 
which have a fibrous structure, and are quite tough and difficult 
to pulverize. It dissolves in two and a half times its weight of 
cold water, and in much less hot water. Its taste is not unpleas- 
antly salty and sharp. Unless in large doses, it is not poisonous. 

142. Ammonium Sulphate, (NH 4 ) 2 SO, is manufactured by 
passing into dilute sulphuric acid the ammonia which is disen- 
gaged when gas liquor is heated with lime. It is in white, color- 
less crystals, readily soluble in water, having a sharp taste. It 
may be used for the manufacture of ammonia, and is employed as 
a fertilizer. 

143. Ammonium Sulphydrate, NH 4 .SH. — We have already 
noticed the composition and mode of formation of potassium 
sulphydrate (§ 101). When hydrogen sulphide is passed into 
ammonia-water until the liquid will dissolve no more of the gas, 
ammonium sulphydrate is formed. 

NH3 + HSH = NH±.SH 

It is a colorless liquid, but becomes yellow after it has been for 

some time exposed to the air. Its odor is disgusting, being at 

the same time that of hydrogen sulphide and that of ammonia. 

If it be mixed with a quantity of ammonia-water exactly equal to 

that from which it was prepared, ammonium sulphide is formed. 

NH 4 SH + NH 3 = NH±.S.NH 4 = (NH 4 ) 2 S 

Ammonium sulphide. 

This compound is of much value in the laboratory in detecting 
some metals. It soon undergoes partial decomposition, and its 
color becomes yellow from the presence of dissolved sulphur. 

To a solution of ferrous sulphate we add a few drops of ammo- 
nium sulphide, and a black precipitate of ferrous sulphide is formed. 

NIH.S.NH* + FeSO* = (NH^SO 4 + FeS 

Ammonium sulphide. Ferrous sulphate. Ammonium sulphate. Ferrous sulphide. 

We pour a few drops of the same liquid into a solution of zinc 
sulphate, and white zinc sulphide is precipitated. 

(NH*)2S + ZnSO* = (NH^SO 4 + ZnS 

144. On examining the composition of the ammonium compounds, we see 
that the radical NH 4 has the same combining power as one atom of hydrogen. 



AMMONIUM AMALGAM. — NITROGEN IODIDE. 103 

It is a monatomic radical ; but at the same time we notice that it can replace 
the hydrogen atoms in the acids, and in so doing it forms salts. It is a basic 
radical, and is in this respect exactly opposite to the radicals CIO- and SO 2 , 
which are acid radicals, 

145. Ammonium Amalgam. — We make an amalgam of 
sodium, — that is, a compound of sodium and mercury, — by throw- 
ing on the surface of a little mercury a few small pieces of sodium. 
If these do not at once combine with the mercury, we can readily 
effect the combination by touching them with a drop of water on 
the end of a long glass rod. As little pieces of burning sodium 
are sometimes thrown out, we keep the vessel at a sufficient dis- 
tance from the eyes. We now pour this amalgam into a tall jar 
containing a strong solution of ammonium chloride : at once a very 
curious phenomenon occurs. The mercury begins to swell and 
become pasty ; it rises and floats on the water, and sometimes it 
overflows the jar. On pouring it out and examining it, we find 
that it has become a brilliant, butter-like substance, and very 
light. It was formerly supposed to be free ammonium dissolved 
in mercury — an amalgam of ammonium, but is probably only 
mercury inflated with hydrogen and ammonia. These gases 
gradually escape and only mercury remains. 

146. Nitrogen Iodide. — We have reduced a small quantity of 
iodine to a fine powder, and we throw this into a little ammonia- 
water. Part of it dissolves, and the other part is converted into 
a black powder, which we carefully pour on a small filter placed 
in a funnel. When most of the liquid has drained off, we dis- 
tribute this powder on several pieces of filter-paper, which we set 
aside for the powder to dry. When it is dry, the lightest touch 
causes it to explode with a loud noise, and sometimes it explodes 
spontaneously. In any case the explosion is always accompanied 
by the production of purple vapors of iodine. The black powder 
is nitrogen iodide : there are several such compounds, and their 
composition depends on the exact manner of formation ; we may 
express it by NP. It is formed by a reaction which yields also 
ammonium iodide. 

4NH3 + 3I 2 = 3N1M + NI 3 

Ammonia. Iodine. Ammonium iodide. Nitrogen iodide. 



104 LESSONS IN CHEMISTRY. 

The ammonium iodide is formed with a considerable produc- 
tion of energy ; but the liquid does not become warm, for all this 
energy is transferred to the nitrogen and iodine atoms which com- 
bine to form nitrogen iodide. Where must we seek the energy of 
explosion of nitrogen iodide ? The explosion is only a rearrange- 
ment of the atoms ; a decomposition of the nitrogen iodide ; the 
energy of this decomposition we must consider as the energy of 
formation of nitrogen molecules and iodine molecules, of which the 
atoms then disengage the energy conferred on them by the forma- 
tion of ammonium iodide, and retained in the nitrogen iodide. 

Nitrogen chloride, NCI 3 , and nitrogen bromide, NBr 3 , have analogous com- 
position. They are oily liquids, and dangerous to prepare. The chloride is 
obtained by inverting a vessel filled with chlorine over warm solution of am- 
monium chloride. It collects in drops, which explode with the utmost vio- 
lence upon heating, in direct sunlight, or in contact with certain substances, 
such as turpentine and phosphorus. 



LESSON XVIII. 
OXIDES OF NITROGEN. 

147. Nitrous Oxide, or Nitrogen Monoxide, N 2 0. — In a 

glass flask, on a sand-bath, we heat some ammonium nitrate, a 
white, crystalline substance obtained by neutralizing ammonia 
with nitric acid. Our flask being provided with a delivery-tube, 
we may collect the gas over the pneumatic trough (Fig. 56). 
The ammonium nitrate is entirely decomposed into water and 
nitrous oxide. 

NH*N0 3 = N20 + 2H20 

Ammonium nitrate. Nitrous oxide. 

As the water is converted into steam by the heat required for the 
experiment, when we desire to collect the gas in a gas-bag we 
pass it first through an empty bottle in which the steam may 
condense. 



NITROGEN MONOXIDE. 



105 



148. Nitrous oxide is a colorless gas, having no odor, but a 
sweet taste. Its density is 22 compared to hydrogen, or 1.527 
compared to air. It is liquefied by great pressure, and con- 
siderable quantities are so liquefied in strong iron cylinders, in 




Fig. 56. 



order that the gas may be transported in small bulk for the use 
of dentists. At ordinary temperatures, water dissolves about its 
own volume of nitrous oxide ; for this reason some of the gas 
is always lost when it is collected over water. 

Nitrous oxide is decomposed by heat, two volumes of the gas 
yielding two volumes of nitrogen and one of oxygen. Since the 
gaseous mixture contains a much larger proportion of oxygen 
than does the air, it should support combustion better than the 
air. An experiment will show us that it does ; we put into a 
jar of nitrous oxide gas a taper bearing only a spark of fire, 
and this spark, decomposing the gas surrounding it, sets free 
sufficient oxygen to relight the taper (Fig. 57). In the same 
manner phosphorus and sulphur burn brilliantly in this gas. 

Nitrous oxide is not poisonous ; it may be inhaled for a short 
time without danger, and its inhalation is followed by insensi- 
bility, a condition called anaesthesia. x\dvantage is taken of this 



106 



LESSONS IN CHEMISTRY. 



property of the gas for the performance of short surgical oper- 
ations. The first effects of the inhalation of the gas are often a 
condition of excitement and disposition to gayety ; for this reason 

it has been called laughing-gas. 
149. Nitric Oxide, NO.— 
In a gas-bottle, provided with 
a delivery-tube and funnel- 
tube, we have some copper clip- 
pings and water. Through the 
funnel-tube we pour nitric acid 
until there is a brisk disengage- 
ment of gas. At first this gas 
in the gas-bottle is red, for 
reasons which we shall presently 
learn, but soon it becomes almost 
colorless. We then pass the de- 
livery-tube under water, and col- 
lect the gas in jars filled with 
water (Fig. 58). In the reaction 
which is taking place, the copper 
is replacing the hydrogen of the nitric acid, and every atom of 
copper replaces two atoms of hydrogen. 

2H 




Fig. 57. 



Cu 

Copper. 



2HN0 3 

Nitric acid. 



Cu(N03)2 
Cupric nitrate. 



But in this case the hydrogen is not set free ; it reduces more 
nitric acid, and if we keep our generating bottle cool by placing 
it in cold water, the reduction yields NO and water. As the cop- 
per atoms always set free even numbers of hydrogen atoms, we 
cannot write this reaction 3H + HNO 3 = 2IPO + NO, but 
must write 6H + 2HN0 3 = 4H 2 + 2NO ; and since the six 
atoms of hydrogen must be replaced by three atoms of copper, 
each of which requires two molecules of nitric acid besides the 
two that are reduced, we may write the whole equation 

3Cu + 8HN0 3 = 3Cu(N03)2 + 4H20 + 2NO 

Although this gas contains but one atom of oxygen in its 



NITRIC OXIDE. 



107 




Fig. 58. 



molecule, it was formerly called nitrogen dioxide : it contains 
twice as much oxygen as the monoxide, or nitrous oxide. 

150. Nitric oxide is a colorless gas, of which we must remain 
ignorant of the taste 
and odor, for it forms 
a corrosive gas as soon 
as it is exposed to the 
air. Its density com- 
pared to air is 1.039. 
It is almost insoluble 
in water. It has been 
liquefied by great cold 
and pressure. 

It is decomposed by J§|| 
heat, but not so readily 
as nitrous oxide. For 
this reason, although 
it contains in a given volume twice the proportion of oxygen 
in nitrous oxide, it will not relight a taper bearing a spark : 
it will, however, support the combustion of phosphorus and 
charcoal. 

The most remarkable property of nitric oxide is its affinity 
for oxygen. We uncover a jar filled with the gas, and instantly 
a cloud of red vapor is formed. This is the red gas which was 
formed in the generating bottle when the nitric oxide first 
eliminated came in contact with the air in the bottle. In this 
experiment each molecule of nitric oxide takes an atom of oxygen 
from the air, and the red vapor is the gas NO 2 . We must be 
careful not to inhale this gas, for it is very injurious. 

We pour a few drops of carbon disulphide into a jar of nitric 
oxide. The vapor of this volatile liquid at once mixes with the 
gas, and when we apply a flame, a bright flash of light fills the 
jar as the carbon is burned by the oxygen of the nitric oxide. 
The light produced by this little explosion affords an excellent 
means for causing the direct combination of hydrogen and chlorine 
(§71). 



108 



LESSONS IN CHEMISTRY. 



We pour a little ferrous sulphate solution into a jar of nitric 
oxide ; some of the gas is at once absorbed, and the liquid becomes 
brown. The nitric oxide may be driven out by heating the solu- 
tion, and the pure gas is sometimes prepared in this manner. 

151. In nitric oxide the affinities of the nitrogen atom are not exhausted : we 
have seen that it is still able to combine with an atom of oxygen. It will also 
combine with an atom of chlorine; when one volume of chlorine is mixed 
with two volumes of nitric oxide, the gases unite, forming a compound NOC1. 
When this compound is treated with water, both substances are decomposed, 
yielding hydrochloric acid and nitrous acid, HNO 2 . 

NOC1 + H20 = HNO 2 + HC1 

Nitric oxide may, then, act as a radical, and in its compounds it is called 
nitrosyl, NOC1 is therefore called nitrosyl chloride, and nitrous acid may be 
called nitrosyl hydrate, NO-OH. 



LESSON XIX. 



OXIDES OF NITROGEN (Continued). 

152. Nitrogen Peroxide, NO 2 and N 2 4 .— We may form this 
substance by the direct combination of nitric oxide and pure oxy- 
gen, and we would of course require two volumes of the first and one 
volume of the second. We can prepare it in another manner. We 

heat some dry lead nitrate 
in a small retort placed in 
a sand-bath. The vapors 
given off are conducted 
into a flask surrounded by 
ice (Fig. 59). Because the 
lead nitrate cannot well be 
perfectly dried, we change 
the receiver after a little 
liquid has collected in it, 
and throw away this first 
portion. That which now collects has a yellow color, and if we 
mix a little salt with the ice around the receiver, the liquid will 




NITROGEN PEROXIDE. 109 

freeze to a crystalline mass. The lead nitrate is decomposed into 

lead oxide, oxygen, and nitrogen peroxide. 

Pb(N0 3 ) 2 = PbO + + N 2 0± 

Lead nitrate. Lead oxide. Nitrogen peroxide. 

The solid nitrogen peroxide melts at — 10° to a nearly colorless 
liquid ; this liquid becomes yellow and afterwards orange-colored as 
the temperature rises, and at 15° is red. It boils at 22°, giving 
the red vapor, and the density of this vapor compared to hydrogen 
is 46, showing that the molecular weight of the compound is 92, 
and the molecule must, therefore, contain N 2 0*. However, as the 
temperature rises the density diminishes, and at 140° it is only 
one-half 46 ; after this, the density remains constant, the mole- 
cule has become two molecules, and each of these must contain 
NO 2 . Such decomposition of gases by heat is called dissocia- 
tion : we have already noticed the dissociation of water vapor 
(§ 54) and of nitrous oxide. 

153. Nitrogen peroxide dissolves in water, but in dissolving 

it reacts with the water ; with a small quantity of water it forms 

nitrogen trioxide and nitric acid, while with a larger quantity 

it yields nitric acid and nitric oxide. 

2N20* + H 2 = 2HN0 3 + N 2 3 

Nitrogen peroxide. Nitric acid. Nitrogen trioxide. 

3N 2 0± + 2H 2 = 4HN0 3 + 2N0 

With the alkaline hydroxides it yields nitrates and nitrites. 

N 2 0± + 2NaOII = NaNO 3 + NaNO'-' + H 2 
Sodium hydroxide. Sodium nitrate. Sodium nitrite. 

A similar decomposition really takes place with water, but the 
nitrous acid formed is at once decomposed by the water. 

The red vapors are dangerous to inhale, and the more dangerous 
because they do not give immediate discomfort. They act on the 
delicate membrane of the lungs, and there have been many fatal 
accidents where the gas has been inhaled by workmen repairing 
sulphuric acid chambers (§§ 108, 109). 

154. Nitrogen Trioxide, N 2 3 .— The existence of this compound is rather 
doubtful. A mixture of equal volumes of nitric oxide and nitrogen peroxide, 
upon cooling to a low temperature, condenses to a blue liquid which is sup- 
posed to be N 2 3 . 



110 LESSONS IN CHEMISTRY. 

The corresponding acid HNO 2 , which would result from the addition of a 
molecule of water to this oxide, has not been isolated. 

N'W + H20 = 2HN02 

Nitrous acid. 
The salts of this acid, or nitrites, are very stable. 

155. Nitrogen Pentoxide, N 2 5 . — When dry chlorine gas is passed over 
silver nitrate, heated in a tube to 70°, the silver and chlorine combine to- 
gether, forming silver chloride ; oxygen is given off and a volatile solid 
compound condenses in the cooler part of the tube. This body is nitrogen 
pentoxide, N 2 5 . 

2AgN0 3 + CI 2 = 2AgCl + N 2 0* + 
It is more readily obtained from nitric acid by withdrawing the elements of 
water by means of phosphorus pentoxide. 

2HN0 3 + P 2 0& = 2HP0 3 + N 2 0* 
Nitrogen pentoxide crystallizes in colorless rhombic prisms ) it melts at 30° 
and decomposes above 45° with evolutions of brown fumes. It is liable to 
explode spontaneously. 

156. In studying the other elements we have examined the combining 
powers of their atoms, compared to that of an atom of hydrogen, — that power 
to which the name atomicity or valence (worth) has been given. What is 
the atomicity of the nitrogen atom ? The molecule of nitrous oxide closely 
resembles in structure the molecule of water: replace by two atoms of nitro- 
gen the two atoms of hydrogen of water, and we have a molecule of nitrous 
oxide. The nitrogen atoms here have the same combining power as the 
hydrogen atoms, and we say they are monatomic. In ammonia, however, the 
nitrogen atom itself combines with three atoms of hydrogen ,• it must then be 
worth three hydrogen atoms, and we call it triatomic. But in ammonium 
chloride it is united with four hydrogen atoms and one chlorine atom ; since 
we have agreed that the chlorine atom has the same worth as the hydrogen 
atom, the nitrogen atom in ammonium chloride must be pentatomic. Now 
let us look at the other oxygen compounds of nitrogen : we have seen that in 
the compounds already studied the oxygen atom is diatomic, and indeed we 
shall in time find that there are many reasons for believing that oxygen is al- 
ways diatomic. Then in nitric oxide, NO, the nitrogen atom, which is combined 
with only one oxygen atom, must also be diatomic; but we have seen that this 
compound NO combines directly with a chlorine atom, forming the compound 
nitrosyl chloride, NO-C1 : it combines with a hydroxyl group, which is mon- 
atomic, forming nitrous acid, HO-NO, and in these compounds the nitrogen 
must be triatomic. We must conclude, however, that in NO 2 the nitrogen is 
tetratomic, since it is combined with two atoms of diatomic oxygen, but here 
ao-ain a hydroxyl group will unite with the nitrogen atom, which is then 
pentatomic in nitric acid, HNO 3 . At low temperatures, when the red vapors 
condense, forming molecules of N 2 4 , it seems also that nitrogen is penta- 



ATOMICITY OF NITROGEN. Ill 

tomic, and that in the molecule of N 2 4 two nitrogen atoms, each of which 
is combined with two oxygen atoms, are also combined with each other. 
If now we remember our representations of the combining powers of the 
atoms by short lines, we may see how the atoms in these molecules seem to be 
related, and how the atomicity of nitrogen varies. 





H 






N-O-N 


H-N-H 


N=0 


C1-N=0 


Nitrous oxide. 


Ammonia. 


Nitric oxide. 


Nitrosyl chloride. 



The reaction between nitrosyl chloride and water then becomes a double 
decomposition which we can easily understand. 

0=N-C1 + H-O-H = 0=N-0-H + H-& 

We can understand also how nitrogen peroxide decomposes by the action of 
water, yielding nitric and nitrous acids. 

O^N=0 0=N=0 

i + H-O-H = • TT + 0=N-OH 

0=N=0 O-H 

Nitrogen peroxide. Water. Nitric acid. Nitrous acid. 

These formulae, which are called constitutional or graphic formulae, do not 
in any manner represent the positions in which the atoms are arranged ; they 
are intended to show what atoms are in relations with other atoms in the 
molecule. We must believe that the atoms in a molecule are in continual 
motion, which we may compare to the motions of the planets around the sun, 
and those of the moons around each particular planet. The nature of the 
molecule depends on all of its atoms, just as the nature of a system of planets 
depends on the central sun and all the planets and their satellites ; and just 
as the moon would go with the earth were that planet to be withdrawn 
from the solar system, so do certain groups of atoms enter into the composition 
of molecules, from which they may separate as groups to form part of other 
molecules or systems of atoms. 

Hereafter we shall not be obliged to use the lines to represent the atomicity 
of all the atoms in the molecules of which we study the structure. We know 
that the group hydroxy!, OH, is monatomic, as are also the groups NO and 
NO 2 ; on the other hand, we know that the group SO 2 is diatomic, and we 
can represent our idea that each of these groups exists in the molecule as a 
distinct part of the system by separating it from the rest of the molecule by a 
period. Thus we may represent nitric acid by the formula N0 2 .OH : sulphuric 
acid, by the formula S0 2 .(OH) 2 . 



112 



LESSONS IN CHEMISTRY. 



LESSON XX. 
NITRIC ACID. HNO 3 . 

157. Minute quantities of nitric aeid often exist in the atmos- 
phere, where they are probably formed under the influence of at- 
mospheric electricity on the nitrogen, oxygen, and moisture of the 
air. Wherever organized matters containing nitrogen decompose 
in the presence of porous substances and alkalies, such as potas- 
sium hydroxide, sodium hydroxide, or lime, nitrates are formed. 
The nitric acid and the nitrates of commerce are manufactured 
from nitrates which are found abundantly in some soils, par- 
ticularly in India, Egypt, and Chili : in the latter country are 
large deposits of sodium nitrate. 

158. We may prepare some nitric acid by distilling in a glass 



JEjfl. .-Q-, 




Fig. 60. 



retort a mixture of sodium nitrate and sulphuric acid, and con- 
densing the vapor in a flask surrounded by cold water. On the 
large scale, the operation is conducted in cast-iron retorts (Fig. 
60), and the vapor is condensed in a series of large stoneware 



NITRIC ACID. 113 

bottles which are called bon-bons. As in the decomposition of 
sodium chloride (§ 75), one molecule of sulphuric acid may be 
made to decompose either one or two molecules of either potassium 
or sodium nitrate, forming at the same time either a neutral or 
an acid sulphate, and setting free one or two molecules of nitric 
acid. 

IPSO* + NaNO 3 = NaHSO* + HNO 3 

Sodium nitrate. Sodium acid sulphate. Nitric acid. 

IPSO* + 2NaN0 3 = Na 2 SO* + 2HN0 3 

The proportion required by the last reaction is that employed 
in the arts, as it is more economical. Let us see what that pro- 
portion must be: the molecular weight of sulphuric acid is 

*' + S + °l = 98 ; that of sodium nitrate is * a + ^ + ° 8 = 85. 

2 + 61 + t>4 23+14+48 

Then 98 parts of sulphuric acid, and 85 of sodium nitrate, if 
perfectly pure, would yield H + N + ° 3 = 63 parts of nitric acid. 

J 1+14 + 48 r 

Properties. — Nitric acid is a colorless liquid, but is partially 
decomposed by the prolonged action of light, red vapor being 
formed and communicating a yellow color to the acid in which 
it dissolves. It is very volatile, and its vapor condenses the 
moisture in the air, producing white fumes. Its density is 1.53. 
It freezes at — 49°, and boils at 86° ; while boiling it is par- 
tially decomposed, so that after a time the boiling point rises to 
120.5°, and a more dilute acid distils, having the same strength 
as that left in the retort. It mixes with water in all proportions, 
and the liquid becomes warm during the mixture. 

159. By a red heat, nitric acid is at once decomposed into 
water, red vapor, and oxygen, a decomposition exactly similar to 
that experienced by lead nitrate under the action of heat (§ 152). 
2HN0 3 = H20 + 2N0 2 + 

In a small crucible, or a thin iron dish, we heat some powdered 
charcoal until it becomes barely red hot. We now remove it from 
the fire, and, when the dish has cooled a little, we pour, at arm's 
length, some strong nitric acid on the still hot charcoal. At once 
a vivid combustion takes place ; the oxygen of the decomposed 

8 



114 LESSONS IN CHEMISTRY. 

nitric acid combines with the carbon, and clouds of red vapor are 
given off. 

On the end of a stick about a metre long we tie a test-tube, into 
which we pour some strong nitric acid, and if our nitric acid is 
not the strongest, we add to it about half its volume of sulphuric 
acid, which will strengthen the nitric acid by its affinity for water. 
Then in another iron dish we carefully warm some good oil of 
turpentine until it is nearly boiling. Now we warm our nitric 
acid, and standing at a distance, pour it suddenly into the hot 
turpentine : at once the nitric acid oxidizes the turpentine, and, 
unless the latter has previously become thick by too long exposure 
to the air, it will be inflamed. 

These experiments show us that the oxygen atoms have not 
exhausted their energy in combining with nitrogen. Indeed, we 
have seen in the conversion of sulphur dioxide into sulphuric acid 
that the oxygen of nitric acid is more energetic than in free oxy- 
gen molecules at ordinary temperatures, for we have to heat oxy- 
gen before it will combine with sulphur dioxide. By reason of 
this energy still existing in its oxygen atoms, nitric acid is easily 
reduced ; that is, part or all of its oxygen may be readily removed 
by oxidizable bodies. We have seen how it is reduced by the 
hydrogen of another portion of the acid when copper replaces that 
hydrogen (§ 149) : in this same reaction part of the nitric acid is 
converted into nitrous oxide and even free nitrogen, so that the 
nitric oxide prepared by nitric acid and copper is never per- 
fectly pure. When the reduction by some metals is carried out 
to its full limit, the nitrogen combines with the hydrogen, forming 
ammonia. This occurs in the action of zinc on very dilute nitric 
acid: although zinc nitrate is then formed, no hydrogen is set 
free, for the displaced hydrogen reduces the nitric acid and com- 
bines with the nitrogen : the ammonia formed at once combines 
with some of the nitric acid present, forming ammonium nitrate. 
HNO 3 + 4H2 = 3H 2 + NH 3 

160. When nitric and hydrochloric acids are mixed, a liquid 
called nitro-hydrochloric acid y or aqua regia, is obtained. This 
liquid is capable of dissolving gold and platinum, a power pos- 



NITRIC ACID. 115 

sessed by neither of the separate acids. We put a small piece 
of gold-leaf in a test-tube with nitric acid, and a similar piece 
in another tube with hydrochloric acid. In neither tube is the 
gold affected, but on mixing the liquids both pieces are dissolved. 
Nitro-hydrochloric acid converts the metals into chlorides, the 
hydrogen of the hydrochloric acid reducing the nitric acid, and 
the chlorine combining with the metal. 

2HN0 3 + 2HC1 = 2H 2 + 2N0 2 + CI 2 

161. Nitrates. — When the hydrogen of nitric acid is replaced 
by metal, nitrates are formed. We have already learned that an 
atom of some metals, which we called monatomic metals, is capa- 
ble of replacing one atom of hydrogen, while the atoms of other 
metals (diatomic) are able to replace two. Since a molecule of 
nitric acid contains only one hydrogen atom, an atom of zinc or 
of lead must replace that atom in two molecules of nitric acid, and 
consequently it will be united to two groups, NO 3 . Then, while 
we can express the molecules of potassium and sodium nitrates by 
the formulae KNO 3 and NaNO 3 , we must write the molecules of 
lead and zinc nitrates Pb(N0 3 ) 2 and Zn(N0 3 ) 2 . 

162. Into a test-tube containing some solution of potassium 
nitrate in water, we pour a little solution of ferrous sulphate, and 
then, inclining the tube, some strong sulphuric acid. This last, 
being much heavier than the other liquids, does not mix at once 
with the solution, but at the surface, where the sulphuric acid 
below and the solution of the nitrate touch, a dark ring is formed. 
This is caused by a partial reduction of the nitric acid by the fer- 
rous sulphate, which produces at the same time a dark color with 
the nitric oxide resulting from the reduction. This color disap- 
pears if we heat the tube (§ 150). This is our test for nitric acid 
and nitrates. 



116 LESSONS IN CHEMISTRY. 

LESSON XXL 
NITRATES. 

163. All of the nitrates are soluble in water. Some of them 
form anhydrous crystals ; others require water of crystallization. 
When thrown on hot coals, they decompose, and the oxygen given 
off increases the intensity of the combustion. Salts which so pro- 
mote combustion are said to deflagrate on hot coals. 

164. Sodium Nitrate, NaNO 3 , is found in large quantities in 
Chili and Peru. It forms rhombohedral crystals that are almost 
cubical ; it is very soluble in water. It attracts moisture from 
the air, and this property prevents its use in the manufacture of 
gunpowder (see § 166). It is from sodium nitrate that nitric 
acid and, indirectly, most of the other nitrates are prepared. 

165. Potassium Nitrate, KNO 3 . — This salt is commonly called 
nitre or saltpetre. In some hot countries it forms an efflorescence, 
or white powder, on the surface of the soil, and may be obtained 
by washing the soil with water and evaporating the resulting solu- 
tion. It is generally made by a double decomposition between 
the sodium nitrate from Chili and either potassium chloride or 
potassium carbonate. Boiling solutions of the two substances are 
mixed, and the sodium chloride or carbonate formed, being much 
less soluble in boiling water than the potassium nitrate, may be 
readily separated. 

KCl -f- NaNO 3 = NaCl + KNO* 

Potassium chloride. Sodium nitrate. Sodium chloride. Potassium nitrate. 

K2C0 3 + 2NaN0 3 = Na 2 C0 3 + 2KN0 3 

Potassium carbonate. Sodium nitrate. Sodium carbonate. Potassium nitrate. 

Potassium nitrate forms long, six-sided prisms which have a 
bitter and cooling taste. They dissolve in about five times their 
weight of water at ordinary temperatures, but require less than 
half their weight of boiling water. We can now understand how 
this compound may be separated from sodium chloride, which is 



GUNPOWDER. 117 

about equally soluble in hot and cold water ; for when a boiling satu- 
rated solution of potassium nitrate is cooled to ordinary tempera- 
tures, nine-tenths of the salt separate in crystals, but from a boiling 
saturated solution of common salt very little is deposited on cooling. 

Potassium nitrate deflagrates — that is, increases the activity of 
combustion — when thrown on hot coals. We melt some zinc in 
an iron ladle, and, when it is nearly red hot, we throw in a few 
small pieces of potassium nitrate : the metal takes fire and burns 
into zinc oxide, the oxygen being supplied from the decomposing 
potassium nitrate. 

166. Gunpowder is a mixture of about seventy-five parts of potas- 
sium nitrate, ten of sulphur, and fifteen of charcoal. It is made 
by grinding each substance separately to the finest powder, and 
then mixing them and grinding, after a little water has been added. 
The intimate mixture is then strongly pressed and carefully dried 
in a warm room, after which it is broken into grains and these are 
sifted into various sizes. The grains are polished by friction over 
each other in rotating barrels. This mixture contains all of the 
materials necessary for its own combustion, and the result of the 
explosion may be generally expressed by saying that the sulphur 
combines with the potassium, forming potassium sulphide, while 
the oxygen of the nitre unites with the carbon to form the gases 
carbon monoxide and carbon dioxide, which, together with the 
nitrogen, are set free. The gas occupies a volume very much 
greater than that of the powder which produced it, and this large 
volume is made still larger by the high temperature of the reaction. 
Of course the outside of the grains of powder must burn first, and 
the larger the grains the slower the combustion and the consequent 
production of gas ; but the smaller the grains the more rapidly is 
each burned and the flame carried from one to the other. Hence 
the small quantity of powder used in small-arms is in fine grains 
in order to produce instantly as much force as possible ; but large 
guns would be broken by such sudden strain, and large grains or 
lumps are employed, which are put into the gun in coarse bags. 
In blasting, if it is desired to break the rock in small pieces, a very 
quickly burning powder is used ; but if it is desired to split off 



118 LESSONS IN CHEMISTRY. 

large masses, the effect is accomplished by the more slowly in- 
creasing pressure from a slower powder. 

167. Silver Nitrate, AgNO 3 , is made by dissolving silver in 
nitric acid, and evaporating the solution until it crystallizes. It 
forms colorless plates, soluble in their own weight of water. Its 
color darkens by the action of the organic matter in the air and 
exposure to light. It melts when cautiously heated, and when cast 
into sticks forms the lunar caustic used by surgeons. It is a cor- 
rosive body, and in the presence of moisture destroys the tissues. 
Should any of it by accident be swallowed, common salt is its anti- 
dote ; insoluble silver chloride is then formed, and this is com- 
paratively harmless (§ 75). When silver nitrate is highly heated, 
it leaves a residue of pure silver. 

168. Strontium Nitrate, Sr(N0 3 ) 2 .— This salt is made by 
dissolving the mineral strontianite, which is strontium carbonate, 
in nitric acid, and purifying by several crystallizations. It forms 
colorless crystals, quite soluble in water. 

169. Barium Nitrate, Ba(N0 3 ) 2 , is obtained like the preceding 
salt, but witherite — barium carbonate — is used. It also forms color- 
less crystals, soluble in water, and the solution may be used as a test 
for sulphuric acid (§ 115). 

170. Cupric Nitrate, Cu(N0 3 ) 2 + 3H 2 0, remains in the bottle 
in which we prepare nitric oxide. If we filter and evaporate this 
solution, the salt separates in large blue prisms ; it is very corrosive. 
When strongly heated, it leaves black, cupric oxide. 

171. Mercuric Nitrate, 2Hg(N0 3 ) 2 +H 2 0, separates in large, 
colorless crystals when we cool in ice and salt the solution obtained 
by boiling mercury in a large quantity of nitric acid. Its solution 
is an energetic caustic, and is used in surgery. When dry mer- 
curic nitrate is heated, it decomposes just as the nitrates of lead 
and copper, leaving red mercuric oxide. 

Hg(N0 3 ) 2 = HgO + 2N0 2 + 

172. Lead Nitrate, Pb(N0 3 ) 2 . — This compound, which is one 
of our most soluble lead salts, is made by boiling lead oxide (lith- 
arge) in nitric acid, and evaporating the solution until crystals 
separate. 



PHOSPHORUS. 



119 



PbO + 2HN0 3 = Pb(N03)2 + H20 
It forms colorless, anhydrous crystals, very soluble in boiling water, 
and in seven times their weight of cold water. 



LESSON XXII. 
PHOSPHORUS— HYDROGEN PHOSPHIDE. 

173. Phosphorus, P=31. — The element phosphorus is extracted 
from bones, in which it exists in a compound known as calcium 




Fig. 61. 

phosphate. The bones are first burned, to remove all of the ani- 
mal matters, and, by a process which we will understand better 
when we study the acids of phosphorus, the calcium phosphate 
which remains is converted into calcium metaphosphate. This 
last body is mixed with charcoal and strongly heated in clay retorts, 
and the phosphorus vapor is condensed in vessels containing cold 
water (Fig. 61). In the reaction which takes place, only half of 
the phosphorus is separated from the bone-ash, and there is left in 
the retorts a compound called calcium pyrophosphate, while the 
gas carbon monoxide is disengaged. 



120 LESSONS IN CHEMISTRY. 

4Ca(P03)2 + IOC = 2Ca2p20* + 10CO + P* 

Calcium metaphosphate. Carbon. Calcium pyrophosphate. Carbon monoxide. 

A simpler process has been recently introduced. It consists in heating a 
mixture of bone-ash, charcoal, and a flux by means of a powerful electrical 
current : phosphorus distils over, while the residue forms a liquid slag. 

Small particles of charcoal are carried over with the phos- 
phorus, which is purified by enclosing it in chamois-skin bags and 
melting it under warm water. The melted phosphorus is then 
squeezed through the leather, and so purified is drawn up into 
glass tubes, where it is allowed to harden in the form of sticks. It 
is always kept under water, and is transported in sealed tin cans. 

174. Properties. — Phosphorus is an almost colorless, wax- 
like solid. It is flexible, and soft enough to be readily scratched 
by the finger-nail. When it has been exposed to light for a long 
time, its surface becomes white and opaque ; it is covered with 
little crystals of phosphorus ; these become loosened, and if we 
shake the bottle in the dark the whole of the liquid is luminous. 
Phosphorus has a peculiar, somewhat garlicky odor. Its density 
is 1.83. It melts at 44°, and boils at 290°. Its density com- 
pared to hydrogen is 62 ; the densities of its gaseous compounds, 
and the composition of all its compounds, show that its atomic 
weight is 31 ; therefore the molecule of phosphorus vapor must 
contain four atoms, if the molecule of hydrogen contains two. 
That is, two volumes of hydrogen weighing 2 represent two atoms, 
but two volumes of phosphorus vapor weighing 124 contain four 
atoms. A number of other elements also have molecules contain- 
ing four atoms. Phosphorus is insoluble in water, but dissolves 
slightly in most oils : it dissolves freely in carbon disulphide, and 
separates in small crystals when the solution is evaporated very 
slowly. Phosphorus is luminous in the dark, and this phenomenon 
is probably caused by a slow oxidation. 

175. Phosphorus has an energetic affinity for oxygen. If we 
expose to the air a small piece of dry phosphorus on a plate, after 
a time the heat developed by the slow combustion is sufficient to 
ignite the phosphorus, which takes fire at a temperature of 50°. 
If we pour on a piece of dry paper on a plate a few cubic centi- 



AMORPHOUS PHOSPHORUS. 121 

metres of a solution of phosphorus in carbon disulphide, the latter 
evaporates, leaving the phosphorus in a state of fine division. 
These small particles are surrounded by oxygen, and the tem- 
perature quickly rises till they burst into flame. 

Phosphorus is very poisonous : even when poisoning by it is not 
rapidly followed by death, dangerous diseases of the liver, heart, 
kidneys, and tongue are produced, and these are usually fatal. 

176. Amorphous Phosphorus. — The properties which have 
just been described are those of the common form of phosphorus, 
but there is another form which may be obtained by heating ordi- 
nary phosphorus for a long time to 240°. It then becomes 
brownish red, opaque, and amorphous. It is not luminous in the 
dark, does not melt at 44° nor take fire at 50°, is insoluble in 
carbon disulphide, and is not poisonous. We may easily make a 
little of this red phosphorus. We put a piece of dry phosphorus 
in a test-tube, and drop on it a very small flake of iodine : the 
iodine combines violently with part of the phosphorus, producing 
light and heat ; but the remainder of the phosphorus has become 
a hard black mass, to extract which we must probably break the 
tube. This black substance is amorphous phosphorus, and when 
powdered is brown. 

While amorphous phosphorus does not take fire as readily as 
ordinary phosphorus, its chemical properties are unchanged. We 
mix a small quantity of moist amorphous phosphorus with pow- 
dered potassium chlorate, and distribute the mixture on several 
pieces of paper which we set aside to dry. When quite dry, the 
least pressure on the spot containing the mixture will cause the 
oxidation of the phosphorus and decomposition of the potassium 
chlorate with a loud explosion. 

When amorphous phosphorus is heated to 260°, it again 
changes into ordinary phosphorus. 

177. Large quantities of phosphorus are employed for the 
manufacture of matches. The flame of phosphorus alone would 
not ignite the stick, because this would become coated with the 
phosphoric oxide formed, and the latter is a bad conductor of 
heat. Common matches are therefore first tipped with paraffin or 



122 LESSONS IN CHEMISTRY. 

sulphur, which may take fire from the phosphorus, and the ends 
of the sticks are then dipped in a paste of ordinary phosphorus 
with strong glue and some coloring matter. The brown-headed 
or parlor matches are tipped with a paste made of amorphous 
phosphorus and potassium chlorate, and sometimes antimony sul- 
phide. The safety matches, which light only on the box, contain 
the potassium chlorate and antimony sulphide, and these are 
ignited by friction with amorphous phosphorus glued to the side 
of the box. 

Burns by phosphorus are quite painful and difficult to heal. 
They are really poisoned wounds, for part of the metaphosphoric 
acid (§ 187), formed by the action of the phosphoric oxide on 
the skin, is absorbed, and the gravity of the burn is much greater 
than that of an ordinary burn of the same size. Phosphorus 
should always be cut under water, and removed from the water 
and dried between folds of filter-paper only at the instant before 
using. 

178. Hydrogen Phosphide, PH 3 . — In a small glass retort 
which we have completely filled with a rather strong solution of 

•sodium hydrate, we put some 
small pieces of phosphorus, 
and after arranging the beak 
of the retort under the sur- 
face of water contained in a 
small vessel, we apply a gentle 
heat (Fig. 62). When the 
liquid begins to boil, bubbles 
_, of gas rise through the water, 

and as each bubble comes 
into the air it takes fire and produces a ring of white smoke. 
When the air is perfectly still, we notice the curious motions of the 
rings. The gas which is being formed is hydrogen phosphide, 
having the composition PH 3 , and as it burns the hydrogen is 
converted into water, and the phosphorus into phosphoric oxide 
which forms the wreaths of smoke. Hydrogen phosphide is not, 
however, the only product of the reaction. Part of the phos- 




OXIDES AND ACIDS OF PHOSPHORUS. 123 

phorus has been oxidized at the expense of some decomposed 

water, and the sodium has entered into the new molecule. As we 

know by analysis that only sodium hypophosphite, having the 

composition NaH 2 P0 2 , and hydrogen phosphide are formed, we 

may write the rather difficult reaction, 

3NaOH + P* + 3H20 = 3XaH2P0 2 + PH* 

Sodium hydroxide. Sodium hypophosphite. 

We must notice that the molecule of hydrogen phosphide has 
a composition like that of ammonia, NH 3 . Indeed, it will under 
proper conditions combine directly with acids, like ammonia, and 
its compounds, which then contain the group PH 4 , are called 
phosphonium salts. 

If we heat red hot in an earthen crucible some fragments of quick-lime, and, 
haviDg a cover for one crucible, throw in some pieces of phosphorus, covering 
the crucible after introducing each piece, a calcium phosphide is formed in 
the crucible. When the crucible and contents have cooled, we may throw some 
of the pieces of the calcium phosphide into water ; bubbles of hydrogen phos- 
phide then come to the surface and take fire spontaneously, forming wreaths of 
smoke as before. Pure hydrogen phosphide does not take fire on coming into 
the air, unless the water through which it passes is boiling. That which we 
have just prepared contains a trace of another compound of phosphorus and 
hydrogen which is spontaneously inflammable. 

179. Phosphorus Chlorides.— There are two chlorides of phosphorus. 
Phosphorus trichloride, PCI 3 , is a volatile, colorless liquid. It is made by pass- 
ing chlorine over phosphorus and condensing the vapor which distils. Phos- 
phorus pentachloride, PCI 5 , is a pale yellow, crystalline solid. It is obtained 
by passing chlorine into the trichloride until the whole becomes solid. Both 
of these bodies are decomposed by water, as we shall presently see. 



LESSON XXIII. 
OXIDES AND ACIDS OF PHOSPHORUS. 

180. There are two oxides of phosphorus, a trioxide, P*0 6 , and 
a pentoxide, P 2 5 . The trioxide is formed when phosphorus is 
slowly oxidized in dry air. The pentoxide, often called phosphoric 
oxide, results when phosphorus is burned in a full supply of air 



124 LESSONS IN CHEMISTRY. 

or oxygen. We place a piece of phosphorus in a small dish on 
a plate, and, after igniting it, cover the dish with a bell-jar. In a 
short time the phosphoric oxide formed settles on the dish and 
sides of the jar in the' form of a snowy-white powder. When we 
sprinkle some drops of water on this powder, a hissing noise is 
heard ; the water and phosphoric oxide combine, producing much 
heat and an acid of phosphorus. The composition of the acid 
which is thus formed depends on the temperature of the water, 
for one molecule of this same phosphoric oxide is able to react 
with one, two, or three molecules of water, forming three different 

acids. 

P2Q5 + H 2 = 2HP0 3 , Metaphosphoric acid. 

P20 5 + 2H 2 = H 4 P 2 7 , Pyrophosphoric acid. 

P 2 5 + 3H 2 = 2H 3 P0 4 , Orthophosphoric acid. 

We shall presently study these acids. Besides these there are 
two others. Of one we have seen the formation of a salt, sodium 
hypophosphite ; the corresponding acid is of course hypophos- 
phorous acid, H 3 P0 2 . The other is formed by the reaction of 
phosphorus trioxide on water, and one molecule of the trioxide 
reacts with six molecules of water, forming four molecules of 
phosphorous acid. 

We then have a series of acids. 

H 3 P0 2 , Hypophosphorous acid. 

H 3 P0 3 , Phosphorous acid. 

H 3 PO, Orthophosphoric acid. 

H 4 P 2 7 , Pyrophosphoric acid. 

HPO 3 , Metaphosphoric acid. 

181. Hypophosphorous Acid may be made by boiling phos- 
phorus with barium hydroxide, Ba(OH) 2 , and by the cautious ad- 
dition of sulphuric acid exactly precipitating the barium from the 
barium hypophosphite formed. After filtering, the liquid is con- 
centrated until a thick syrup is obtained. This is hypophosphorous 
acid. Although a molecule of this acid contains three atoms of 
hydrogen, only one of those atoms is replaceable by metal. It is 
a monobasic acid, and its salts with a monatomic metal like sodium 
will contain one atom of metal and the group H 2 P0 2 . The hypo- 
phosphites of diatomic metals must contain two of these groups 



ORTHOPHOSPHORIC ACID. 125 

in order that two atoms of hydrogen may be replaced : barium 
hypophosphite will, then, be Ba(H 2 P0 2 ) 2 . 

182. Phosphorous Acid is most quickly prepared by the re- 
action of phosphorus trichloride with water, one molecule of the 
trichloride requiring three molecules of water. 

PCI 3 + 3H 2 = H 3 P0 3 + 3HC1 
It is a dibasic acid : it contains two atoms of replaceable hydro- 
gen; we may have a sodium phosphite, Na 2 HP0 3 , and a sodium 
acid phosphite, NaH 2 P0 3 . Barium phosphite would be BaHPO 3 . 
Both hypophosphorous and phosphorous acids have reducing 
properties; that is, they will take away oxygen from oxidized 
bodies, so becoming converted into phosphoric acid. Into a test- 
tube containing a solution of silver nitrate, we pour some solution 
of sodium hypophosphite : in a short time the interior of the tube 
is coated with metallic silver by the reducing action of the hypo- 
phosphite. 

183. Orthophosphoric Acid. — We have seen how this acid, 
which is commonly called phosphoric acid, may result from the 
action of water on phosphoric oxide. It is also formed by the 
reaction of phosphorus pentachloride with water. 

PCI 5 + 4H 2 = H 3 PO* + 5HC1 
It is usually made by boiling amorphous phosphorus with nitric 
acid, which is reduced, red vapors being given off. The liquid is 
then evaporated to a small bulk, and put in a bell-jar over a dish 
containing sulphuric acid, which gradually absorbs the remain- 
ing moisture. In this manner hard, transparent, and deliquescent 
crystals of orthophosphoric acid are obtained. 

Orthophosphoric acid is tribasic : its molecule contains three 
atoms of replaceable hydrogen. Consequently it may with the 
same metal form three different salts, accordingly as one, two, or 
three atoms of hydrogen are replaced by a corresponding quantity 
of the metal. The names of these salts should indicate the number 
of hydrogen atoms which have been replaced, or the number of 
metallic atoms which have replaced the hydrogen : thus, since one 
atom of sodium always replaces one of hydrogen, monosodium 
phosphate is NaH 2 P0 4 , disodium phosphate is Na 2 HP0 4 , and tri- 



126 LESSONS IN CHEMISTRY. 

sodium phosphate is Na 3 P0 4 . We have already learned by several 
reactions (§§ 118, 136) that one atom of calcium is capable of re- 
placing two atoms of hydrogen ; and if we perfectly neutralize 
orthophosphoric ^cid with lime (calcium oxide), we must have 
two molecules of the acid and three of lime. 

2H 3 P0 4 + 3CaO = Ca 3 (P0 4 ) 2 + 3H 2 
The tricalcium phosphate so formed is the compound existing 
in bone-ash, from which phosphorus is obtained. It is insoluble 
in water ; when it is treated with sulphuric acid, two atoms of 
calcium are taken from its molecule, forming calcium sulphate, 
while calcium acid phosphate passes into the solution. 

Ca 3 (PO±) 2 + 2H 2 S0* = CaH*P0 4 + 2CaS0* 

Tricalcium phosphate. Calcium acid phosphate. Calcium sulphate. 

The calcium sulphate, being insoluble, is separated by filtration, 

and the calcium acid phosphate is converted into calcium meta- 

phosphate by the action of heat, which decomposes it with the 

formation of water. 

CaH*(P0 4 ) 2 = Ca(P0 3 ) 2 + 2H 2 

Calcium acid phosphate. Calcium metaphosphate. 

184. To a solution of disodium phosphate — either of the other 

orthophosphates would answer — we add a little ammonia-water, 

and then some magnesium sulphate solution. A white precipitate 

forms ; this contains both ammonium and magnesium ; two atoms 

of hydrogen in phosphoric acid are here replaced by one atom of 

magnesium, and the other by the ammonium group } NH*. 

Na 2 HPO* + MgSO + NH 3 = Na 2 SO* + Mg(NH 4 )PO* 
Disodium Magnesium Sodium Ammonio- 

phosphate. sulphate. sulphate. magnesium phosphate. 

In another test-tube we mix some solutions of disodium phos- 
phate and silver nitrate. A yellow precipitate of trisilver phos- 
phate forms. 

Na 2 HP0 4 + 3AgN0 3 = Ag 3 P0 4 + 2NaN0 3 + HNO 3 

These reactions enable us to identify orthophosphoric acid and 
the orthophosphates. 

185. Orthophosphates. — Disodium phosphate exists in the 
blood, and the phosphorus which is eliminated from our bodies is 



PYROPHOSPHORIC ACID. 127 

principally in monosodium phosphate, which passes out in the 
urine. The phosphates containing only one atom of metal redden 
blue litmus, and are generally called acid phosphates. Those 
containing two atoms of metal do not affect litmus, and are gen- 
erally called neutral or common phosphates ; while those having 
three atoms of metal turn red litmus to blue. 

Large mineral deposits of tricalcium phosphate exist in many 
localities, and it is probable that they have been formed from ac- 
cumulations of bones during prehistoric ages. The mineral apa- 
tite, generally green in color, is principally tricalcium phosphate. 

186. Pyrophosphoric Acid. — When orthophosphoric acid is 
long heated to a temperature of about 213°, it undergoes partial 
decomposition : two molecules lose one molecule of water, and then 
combine together, forming a molecule of pyrophosphoric acid. 

2H 3 PO± = H20 + H*P 2 7 

We can understand this better if we consider the structure of the molecule 
of phosphoric acid : it must contain three hydroxy! groups, and the other 
atom of oxygen must be combined directly with the phosphorus atom. By the 
removal of the elements of one molecule of water, two groups, each containing 
one phosphorus atom, one oxygen atom, and two hydroxyl groups, will be 
cemented, we may say, by an atom of oxygen. 

OH OH OH OH 

HO-PrO + HO-P^O = OrP-O-PrO + HOH 
OH OH OH OH 

Two molecules orthophosphoric acid. Pyrophosphoric acid. 

We see then that in certain compounds, such as hydrogen phosphide and 
phosphorus trichloride, phosphorus is triatomic, but that in other cases, and 
these are the most numerous, it is pentatomic, or equivalent to five atoms of 
hydrogen. 

We mix some solutions of sodium pyrophosphate and silver 
nitrate ; instantly a white precipitate of insoluble silver pyrophos- 
phate is formed. 

Na*PW + 4AgN0 3 = Ag*P 2 7 + 4NaN0 3 

187. Metaphosphoric Acid. — When either orthophosphoric 
or pyrophosphoric acid is heated to redness, water is formed, and 
there remains a hard, glass-like mass of metaphosphoric acid. 

H 3 P0 4 = HPO 3 - H2Q 



128 LESSONS IN CHEMISTRY. 

If an acid phosphate is heated in the same manner, it undergoes 
a similar decomposition, and a metaphosphate remains (§ 183). 

Metaphosphoric acid quickly coagulates or renders insoluble the 
albumen of white of egg, a property which distinguishes it from 
both ortho- and pyrophosphoric acids. 

Metaphosphoric and pyrophosphoric acids and their salts are 
poisonous, as are also hypophosphorous and phosphorous acids, but 
orthophosphoric acid and the orthophosphates are not poisonous 
unless in such concentrated form as to be corrosive. 

188. When either metaphosphoric or pyrophosphoric acid, or 
any of their salts, is boiled with nitric acid, orthophosphoric acid or 
one of its salts is formed. We have already seen that phosphorus 
itself is oxidized to orthophosphoric acid by nitric acid. If 
to this solution in nitric acid we add a solution of ammonium 
molybdate also in nitric acid, at once or after a time a bright- 
yellow precipitate of a body called ammonium phosphomolybdate 
separates. In this manner we can detect the presence of phos- 
phorus or any of its componnds. 



LESSON XXIV. 
ARSENIC. As = 75. 

189. Arsenic is found associated with many metals, copper, 
silver, bismuth, nickel, but it is obtained principally from one of 
its minerals, which contains also iron and sulphur. This mineral 
is called mispickel, and its composition may be represented by the 
formula FeSAs. When it is strongly heated, the arsenic is driven 
out, and iron sulphide, FeS, remains. The operation is conducted 
in clay retorts, and the arsenic condenses in sheet-iron receivers. 
This impure arsenic is generally sold under the name cobalt ; it is 
purified by being redistilled out of contact with air. 

190. In a small test-tube we heat some commercial arsenic, and 
soon a bright steel-gray ring forms in the cooler part of the tube ; 
after a time the interior of the ring becomes lined with small but 



1 



ARSENIC. 129 

brilliant metallic crystals. This is the appearance of arsenic, but 
its surface oxidizes after some exposure to the air, and becomes 
tarnished. The density of arsenic is 5.7. It does not melt when 
heated ; it sublimes ; but it may be melted to a transparent liquid 
'by heating it under pressure. It is insoluble in water, but is 
slowly oxidized by the air dissolved in the water, and the oxide 
dissolves, rendering the water poisonous. When arsenic is heated 
in contact with air, it volatilizes, and its vapor is oxidized to white 
arsenious oxide. Arsenic takes fire spontaneously in chlorine, 
burning into arsenic chloride, AsCl 3 , which is a volatile, very poi- 
sonous liquid. 

A small quantity of arsenic is added to the lead for making 
shot ; it hardens the shot, and the interior of the gun-barrel does 
not become coated with lead by friction with that soft metal. 

191. Arsenious Oxide, As 4 6 . — We heat a very small frag- 
ment of arsenic in a test-tube, and presently a white ring con- 
denses on the sides of the tube. The arsenic has been oxidized, 
and the volatile oxide has condensed in the tube : if we examine 
the ring by the aid of a good microscope, we find that it is com- 
posed of small, eight-sided crystals. These are arsenious oxide. 
Arsenious oxide is manufactured in this manner, by heating 
arsenic in contact with the air, and the vapor is condensed either 
in cool chimneys or in large rooms. As it is a very poisonous 
substance, the operation is conducted with all possible precaution 
that the workmen may not inhale the vapors and dust. 

When it is freshly sublimed in large masses, arsenious oxide is 
a glassy, transparent, and amorphous solid, but it soon becomes 
opaque, and this is due to the formation of little crystals. It is 
not very soluble in water, and the amorphous form is more sol- 
uble than the crystalline or opaque variety. Amorphous arsenious 
oxide dissolves in twenty-five times its weight of cold water, but 
the crystalline form requires eighty times its weight. The solution 
contains arsenious acid, but when we evaporate the liquid and try 
to separate this acid, it is again decomposed into arsenious oxide 

and water. 

As*0 6 + 6H20 = 4H3AsO s 

Arsenious oxide. Water. Arsenious acid. 

9 



130 LESSONS IN CHEMISTRY. 

Because arsenious oxide is frequently the cause of intentional 
or accidental poisoning, it is important that we shall be able to 
recognize it; but we will better understand its tests when we have 
learned something of the other compounds of arsenic. 

192. Arsenic Oxide and Acids. — When arsenic or arsenious 
acid is boiled with nitric acid, it is oxidized just as phosphorus 
was oxidized, and the ortho-arsenic acid formed corresponds ex- 
actly to orthophosphoric acid. It contains H 3 As0 4 . When it is 
heated to 150°, it is decomposed like orthophosphoric acid, and 
pyroarsenic acid, H 4 As 2 7 , is formed. This also is decomposed 
at 200°, yielding metarsenic acid, HAsO 3 , which when heated to 
redness loses the elements of water, and leaves arsenic oxide, As 2 5 . 

2H 3 AsO± = HUsW + H 2 H±As 2 07 = 2HAs0 3 + IPO 
2HAs0 3 = As 2 5 + H 2 

193. We boil a few grains of arsenious oxide with a few drops 
of nitric acid in a test-tube, and when the last particle of the 
solid disappears, we carefully neutralize the liquid with ammonia. 
Now when we add some silver nitrate solution, a brick-red pre- 
cipitate of silver arsenate is formed. 

(NH*) 3 AsO* + 3AgN0 3 = Ag 3 AsO* + 3NH*N0 3 

Ammonium arsenate. Silver nitrate. Silver arsenate. Ammonium nitrate. 

194. Arsenic Sulphides. — In a test-tube of hard glass we 
melt together some powdered arsenic mixed with a little more 
than half its weight of sulphur. After cooling, the liquid solidi- 
fies to a red mass of arsenic disulphide, As 2 S 2 . This substance 
is commonly called realgar. It is found as a mineral in trans- 
parent red prisms. It is insoluble in water. When heated in 
the air, both its arsenic and sulphur burn, yielding arsenious oxide 
and sulphur dioxide. 

In another test-tube we melt a mixture of powdered arsenic 
with about two-thirds its weight of sulphur. When this tube 
cools, we find in it yellow arsenic trisulphide, As 2 S 3 , generally 
called orpiment. This sulphide also is found as a mineral. It is 
insoluble in water, but if boiled for a long time with that liquid 
it is decomposed, yielding hydrogen sulphide and arsenious acid. 
As 2 S 3 + 6H 2 - 2H 3 As0 3 + 3H 2 S 



TESTS FOR ARSENIC. 131 

Conversely, by passing hydrogen sulphide through a solution 
of arsenious oxide to which a drop of hydrochloric acid has been 
added, yellow arsenious sulphide is precipitated. 

195. Tests for Arsenic. — Arsenious oxide and some of its 
compounds are the usual forms in which we must identify arsenic. 
In a porcelain evaporating dish we heat some pure water, and 
when it boils we add a few drops of hydrochloric acid, and then 
put in a thin strip of bright copper. The metal does not tarnish ; 
but when we add to the boiling liquid a little of any solution con- 
taining arsenic, the copper soon becomes coated with a steel-gray 
or even black deposit, which is a compound of copper and arsenic. 
This is called Reinsch's test. We take this slip of copper from 
the liquid, wash it 

in pure water, and 
carefully dry it be- 
tween folds of warm 
filter-paper. Then 
we cut it into sev- 
eral very narrow 
slips, and put one 
or two of these in a 
little tube drawn out 

and sealed at one 

j ttt i • FlG - 63 - 

end. We cover this 

piece of copper with some warm charcoal powder, and then heat 
the end of the tube. The heat drives the arsenic away from the 
copper, and the charcoal prevents the vapor from becoming oxi- 
dized, so that a gray or black mirror of arsenic condenses in the 
nearest cool part of the tube (Fig. 63). 

196. In another similar tube we put another piece of our coated 
copper foil, and heat it alone. In this case the arsenic vapor be- 
comes oxidized by the air in the tube, and white arsenious oxide 
is deposited in minute octahedral crystals that we may recognize 
when we examine the tube under the microscope (Fig. 64). Were 
we to break off the portion of the tube containing this deposit, 
and boil it with a very little water in another tube, we would 




132 LESSONS IN CHEMISTRY. 

obtain a solution of arsenious acid, with which we could make the 
next tests ; these, however, we will make with larger quantities of 
the substance. 

197. To a solution of arsenious acid in a test-tube, we add a 
drop of ammonia, and then some silver nitrate solution. A canary- 
yellow precipitate of insoluble silver arsenite is formed. 

(NH±)2HAs0 3 + 2AgN0 3 = 2NH 4 N0 3 + Ag2HAs0 3 
Ammonium arsenite. Silver arsenite. 

198. To a similar solution, treated with a little ammonia, we 
add cupric sulphate dissolved in water. 
An apple-green precipitate of cupric 
arsenite, CuHAsO 3 , is thrown down. 

199. In another tube we acidulate 
some arsenious acid with a drop of hy- 
drochloric acid, and then pass hydrogen 
sulphide through the liquid. A bright 
yellow precipitate of arsenic trisulphide 
is formed. 

200. When arsenious acid is poured 
into a bottle in which hydrogen is being generated, the nascent 
hydrogen, that is, the free atoms of hydrogen which have not 
exhausted part of their energy by combining to form molecules, 
will reduce or take away oxygen from the arsenious acid, and 
combine with the arsenic, forming an exceedingly poisonous gas, 
of which we must be careful not to inhale the least quantity. It 
is called hydrogen arsenide. Its molecule contains AsH 3 . 

H 3 As0 3 + 6H = 3H20 + AsH* 
We have prepared a hydrogen-bottle with a long jet (Fig. 65), 
and, while the hydrogen is burning with its pale flame at this jet, 
we pour through the funnel-tube a few drops of a solution of 
arsenious oxide. In a few moments the flame becomes bluish and 
elongated. Hydrogen arsenide is burning, and the arsenic oxidizes 
to arsenious oxide, producing a white smoke. In this flame, and 
close to the jet, we hold a plate or piece of cold porcelain, which 
will prevent the arsenic from getting enough oxygen to become 
oxidized. We see a dark spot of arsenic forming, and we make 







TESTS FOR ARSENIC. 



133 




Fig 



several of these spots on different portions of the plate. This is 

called Marsh's test. If with a lamp we heat the tube of the long 

jet, the hydrogen arsenide 

will be decomposed by 

the heat, and the dark 

ring of arsenic deposited 

in the cooler part of the 

tube may be afterwards 

tested as we have already 

studied. 

201. We connect a 
little bent tube with our 
jet, and pass the gas into 
some silver nitrate solution 
in a test-tube ; a black de- 
posit of silver separates, 
and arsenious acid is 
formed in the solution. 

AsH 3 + 6AgX0 3 + 3H 2 = H 3 As0 3 + 6HN0 3 + Ag6 
We filter the liquid from the silver, and add a drop of ammonia ; 
if all of the silver nitrate has not been decomposed, a yellow 
precipitate of silver arsenite is formed. We may be obliged to add 
a few more drops of silver nitrate (§ 197). 

202. Now we touch one of the spots on our plate with a drop 
of strong nitric acid. The spot disappears : we add a small drop 
of ammonia, and cautiously warm the plate until it is dry. Then 
we touch it with a drop of silver nitrate, and it becomes brick-red 
from the formation of silver arsenate (§ 193). 

All of these tests enable us to recognize arsenic with certainty; 
they are applied to substances which are extracted from the body 
in cases of supposed poisoning. 

The green coloring matters known as Scheele's green and Paris 
green are compounds containing arsenic and copper: they are 
exceedingly poisonous. 



134 LESSONS IN CHEMISTRY. 



LESSON XXV. 

ANTIMONY. Sb (Stibium) = 120. 

203. Antimony is found principally in combination with 
sulphur in a grayish-black mineral, stibnite, Sb 2 S 3 . This sul- 
phide is quite fusible, and it is separated from the earthy matters 
with which it is mixed, and which are called the gangue, simply 
by heating the ore ; the antimony sulphide melts and runs out. The 
easiest method of obtaining antimony from this sulphide is to mix 
the powdered sulphide with scrap iron and heat the mixture to 
redness in a crucible. Iron sulphide and antimony are formed, 
and the latter, being the heavier, collects at the bottom, where we 
find it as a bright button-shaped lump when we break the cold 
crucible. The cheapest method, however, is to roast the powdered 
sulphide ; that is, heat it in the air ; most of the sulphur is then 
oxidized to sulphur dioxide, which passes off, and most of the anti- 
mony is converted into antimonous oxide. The roasted mass is 
then mixed with charcoal, and the mixture moistened with a solu- 
tion of sodium hydroxide, after which it is heated in crucibles. 
The carbon removes the oxygen from the antimony oxide, and the 
sodium hydroxide removes the sulphur from the antimony sulphide 
still present. The sodium sulphide produced forms a slag which 
floats on the surface of the melted antimony. 

Properties. — Antimony is a very brilliant, white substance, 
having a high metallic lustre. It is very brittle, and breaks 
in shining layers : it is said to have a laminated structure. Its 
density is 6.7. It melts at 450° ; when a considerable quantity 
of it is melted in a crucible and allowed to cool quietly until 
a crust forms on the surface, if we make a hole in this crust and 
pour out the still molten interior, the crucible will be found to be 
lined with small shining crystals. 

When antimony is heated in the air, it is oxidized to antimo- 
nous oxide, Sb 4 6 . We have already seen that antimony burns 



ANTIMONY. 135 

spontaneously when thrown into chlorine. It combines with 
the chlorine, forming a trichloride, SbCl 3 , and a pentachloride, 
SbCl 5 . 

Antimony enters into the composition of several alloys. Type- 
metal contains twenty per cent, of antimony and eighty per cent, 
of lead. Lead is too soft for type, and it does not take sharp im- 
pressions of moulds : the antimony renders the metal hard, and 
causes it to expand on solidifying, so filling every line of the 
mould. Britannia metal also contains antimony. 

204. Antimony Chlorides. — By distilling antimony trisulphide with hydro- 
chloric acid, and collecting apart the product which passes after the condensed 
liquid begins to crystallize in the neck of the retort, antimony trichloride, SbCl 3 , 
is obtained as a transparent, colorless solid, melting at 73°, and boiling at 230°. 
It is soluble in dilute hydrochloric acid, but when the solution is diluted with 
water, an insoluble oxychloride, SbOCl, is thrown down, while hydrochloric 
acid is formed. Antimony pentachloride, SbCl 5 , is a volatile, yellow liquid, 
formed by the action of an excess of chlorine on antimony or the trichloride. 

205. Antimony Oxides. — Antimonoits oxide, Sb 4 6 , is made by heating anti- 
mony to redness in open crucibles ; after cooling, the latter are found lined with 
shining, needle-like crystals of the oxide, which corresponds in composition to 
arsenious oxide. When antimony is boiled with strong nitric acid, it is con- 
verted into metantimonic acid, HSbO 3 ; by heating to about 275° this is de- 
composed, yielding antimony pentoxide, Sb 2 5 ; at higher temperatures it 
breaks up into the tetroxide, Sb 2 4 , and oxygen. There is also a pyranti- 
monic acid, H 4 Sb 2 7 , but there is no orthantimonic acid of the composition 
H3SbO*. 

Antimony tetroxide is also formed by heating antimonous oxide in the air. 
It is white powder and insoluble in water. 

206. Antimony Trisulphide, Sb 2 S 3 , is the grayish-black min- 
eral, called stibnite, from which we have already learned that 
antimony is obtained. It is a heavy, crystalline substance, 
having a marked metallic appearance. This same sulphide 
may be obtained in another form. Through a solution of anti- 
mony trichloride we pass hydrogen sulphide ; an amorphous, 
orange-colored precipitate is formed, and this is antimony trisul- 
phide. 

2SbC13 + 3H 2 S = Sb 2 S3 + 6HC1 

Antimony trichloride. Antimony trisulphide. 

207. When a solution containing antimony is introduced into 
a bottle in which hydrogen is being generated, some of the hydro- 



136 LESSONS IN CHEMISTRY. 

gen combines with the antimony, producing a gas, hydrogen anti- 
monide. Although this gas has not been obtained in a pure state, 
being very easily decomposed, enough has been learned about it 
to show that it has the composition SbH 3 . It causes the hydro- 
gen to burn with a bluish flame, somewhat like that of hydrogen 
arsenide, and it also produces dark spots on a piece of porcelain 
held in the flame, as well as rings in the heated tube (§ 200) ; but 
here the resemblance with arsenic ceases. When the spots are 
oxidized by nitric acid and then treated with silver nitrate, no 
brick-red color is produced. When the gas is passed through 
silver nitrate solution, a dark compound of silver and antimony is 
precipitated, and, as the clear liquid then contains only nitric acid, 
it cannot give a precipitate when neutralized with ammonia 
(§§ 200-202). 

208. When we compare the compounds of nitrogen, phosphorus, arsenic, 
and antimony, we find that the atoms of these elements are almost alike as 
far as their power of combining is concerned. One atom of each will combine 
with three atoms of hydrogen, and we then have formed the four gases, 

NH3 PH 3 AsH 3 SbH 3 

Their more important compounds with chlorine show the same similarity : 

NCI 3 PCI 3 AsCl 3 SbCl 3 

But, in addition to these chlorides, phosphorus and antimony form penta- 
chlorides, PCI 5 and SbCl 5 . Each of the four elements has a trioxide and a 
pentoxide, and from «ach of the pentoxides is derived an acid containing one 
atom of hydrogen, one atom of the element, and three atoms of oxygen : 

UNO 3 HPO 3 HAsO 3 HSbO 3 

In addition, phosphorus and arsenic form the ortho-acids, H 3 PO* and 
H 3 As0 4 , while phosphorus, arsenic, and antimony form the pyro-acids, H 4 P 2 7 , 
H*As20 7 , and H*Sb20 7 . 

These similarities and many others enable us to group together the four ele- 
ments in a natural class; since in their compounds one atom of either of the 
class has a combining power equal to that of three or of five atoms of hydro- 
gen, we may call the class the group of triatomic or pentatomic non-metals. 

209. There are three other elements which would be placed in the class that 
we have just considered, but they occur in such small quantities, although 
widely distributed, that we can only mention their names. They are vana- 
dium, niobium, and tantalum. They are found in the minerals vanadanite, 
columbite, and some others. 



BORON. 137 



LESSON XXVI. 

BORON. B= ii. 

210. The well-known substance, borax, is a compound of the 
element boron. To a saturated solution of borax we add some 
sulphuric acid : soon there separates a deposit composed of small 
white flakes. If we filter these flakes from the liquid, and dry 
them, we have pearly white scales which feel greasy like soap 
when we take them between the fingers. This substance is boric 
acid ; it contains H 3 B0 3 . If we heat it red hot in a platinum 
crucible, it decomposes and leaves boric oxide, B 2 3 . 

2H 3 B0 3 = 3H 2 + B20 3 

When this boric oxide is mixed with pieces of sodium, some 

common salt being added to make the mixture melt more readily, 

and heated to bright redness in a covered iron crucible, sodium 

borate and boron are formed. 

3Na 2 + 2B20 3 = 2Na 3 B0 3 + B 2 
Boric oxide. Sodium borate. 

After the crucible has cooled, the fu^ed mass is treated with dilute 
hydrochloric acid, which dissolves the sodium borate, leaving the 
boron as a dark-brown or olive powder. 

Boron is amorphous, infusible, insoluble in water. When it is 
heated in the air or in oxygen, it takes fire and burns to boric 
oxide. It is one of the few elements which combine directly with 
nitrogen : at a red heat in an atmosphere of nitrogen it is con- 
verted into boron nitride, BN. It also burns in nitrogen dioxide 
when heated in that gas, and forms a mixture of boron nitride and 
boric oxide. 

211. Boric Oxide. — Boric oxide, of which we have already 
learned the manner of formation, is a hard, transparent, glass-like 
substance. It melts at a red heat, and when melted has the prop- 
erty of dissolving many metallic oxides, which communicate vari- 
ous colors to the cooled oxide. We heat to redness the end of a 



138 LESSONS IN CHEMISTRY. 

small platinum wire, and, when it is very hot, we dip it into some 
boric acid or powdered boric oxide ; on again heating this in the 
flame, it melts to a sort of glass bead, which is perfectly trans- 
parent and colorless when cold. We now dip it into a solution 
of cobalt chloride, and again heat it : when it cools, the bead has 
a blue color. This blue color is given by the cobalt. As many 
metals give peculiar colors to such beads of boric oxide, we have 
in that substance a valuable reagent to aid in the detection of the 
metals. 

Boric oxide is not reduced by heating it with charcoal, but 
when chlorine is passed over a red-hot mixture of boric oxide and 
charcoal, carbon monoxide, CO, is disengaged, together with the 
vapor of a very volatile liquid, boron chloride, BCP. 

B 2 3 + 3C + 3C1 2 = 2BC1 3 + 3CO 
Boric oxide. Boron chloride. 

212. When boric oxide is melted with the metal aluminium, a 
part of the metal is oxidized, and another part combines with the 
boron from which the oxygen was removed ; there is so formed a 
complex compound of boron and aluminium, which separates in 
small crystals when the mass cools. As these crystals are mixed 
with the excess of solid aluminium, we must remove that metal 
by boiling in dilute hydrochloric acid. Small octahedral crystals 
remain undissolved : they were long regarded as crystallized boron. 
Their composition is not always the same. Their color is yellow, 
red, or black : their density is about 2.6, and they are almost as 
hard as diamond : they will scratch rubies, and have sometimes 
been employed for polishing precious stones. 

213. Boric Acid and Borates. — We have already seen that 
boric acid may be formed by the action of sulphuric acid on borax. 
It dissolves in about twenty-five times its weight of cold water, and 
the solution is not very strongly acid ; it changes blue litmus to a 
wine color. 

It is found in nature in the craters of some volcanoes. In nu- 
merous localities in Tuscany gases issue from cracks in the earth, 
and these volcanic gases contain boric acid. To obtain this body 






BORIC ACIDS. — BORAX. 139 

the gas is caused to bubble through the water of little lakes, and 
when the water is evaporated, the boric acid is left. 

Boric acid is tribasic : when it is heated to 100°, it is decom- 
posed into water and metaboric acid. 

H3B0 3 = HBO 2 + H20 

Boric acid. Metaboric acid. 

If the latter be heated to 140° for a time, it is further decom- 
posed into another acid, called tetraboric. 

4HB0 2 = H2B*07 + H20 
Metaboric acid. Tetraboric acid. 

Tetraboric acid is that to which correspond borax and the 
common borates. In borax, which is sodium tetraborate, both of 
the hydrogen atoms are replaced by sodium. 

214. Borax, Na 2 B 4 7 , was for a long time obtained princi- 
pally from Asia and from the boric acid of Tuscany, but within 
recent years it has been found in large quantities in certain lakes 
(Borax Lake, Lake Clear) in California. The lakes are dredged, 
and the mud from the bottom is boiled with water ; the borax 
separates in crystals when the solution is evaporated. When a 
very concentrated solution of borax cools, it deposits, between 79° 
and 56°, octahedral crystals in which one molecule of borax is 
combined with five molecules of water of crystallization ; but 
below 56° it deposits prismatic crystals containing ten molecules 
of water. As a given quantity of borax will yield a much heavier 
weight of the latter crystals than of the former, the prismatic 
crystals are those prepared for commerce. Prismatic borax dis- 
solves in twelve times its weight of cold, or twice its weight of 
boiling water. 

When borax is heated, it loses its water of crystallization more 
quickly than that water can evaporate ; the borax is consequently 
dissolved in the separated water : it is said to melt in its water of 
crystallization. As the water is driven off by the heat, it causes 
the borax to swell up, until it becomes a dry, white, and very 
light mass. When this is still further heated, it melts to a sort of 
glass, which possesses the same property of dissolving metallic 
oxides that we noticed in fused boric oxide. For this reason borax 



140 



LESSONS IN CHEMISTRY. 



is often used in analysis instead of boric oxide. Because borax 
dissolves metallic oxides, it is useful in brazing and welding. The 
surfaces of metal to be welded together become oxidized at the 
high temperature necessary, and the oxide would prevent their 
union : a little borax sprinkled on the hot surfaces dissolves the 
oxide, and the liquid is squeezed out when they are pressed 
together, leaving clean surfaces which readily unite. 

215. We dissolve in alcohol a little boric oxide, or some borax 
to which a few drops of sulphuric acid have been added. On 
lighting this alcohol, it burns with a green flame. This test helps 
us to recognize boric acid or a borate. 



LESSON XXVII. 



SILICON. Si 



28. 



216. Silicon is one of the most abundant elements. In the 
form of oxide it exists in quartz, and forms part of nearly all 
rocks and of many minerals. It is obtained in the free state by 
heating a mixture of powdered quartz and magnesium powder. 

SiO 2 + 2Mg = 2Si + 2MgO. 

After washing the residue with dilute hydro- 
chloric acid, in which the magnesium oxide 
dissolves, the silicon remains as an amorphous 
brown powder. Two crystallized modifications 
have been prepared. 

217. Silicic Oxide, SiO 2 .— This compound, 
generally called silica, is found in many forms. 
Crystallized, it constitutes the various kinds of 
quartz, such as rock crystal and amethyst ; 
agate, flint, and chalcedony are other varieties. 
Sandstones and sand are also silica. Pure 
quartz or rock crystal is colorless. It forms hexagonal prisms 
terminated by six-sided pyramids, and the angles are often 




Fig. m. 



SILICA. — GLASS. 141 

curiously modified (Fig. 66). Its density is 2.65. It is in- 
soluble in water ; and can be melted only in the oxyhydro- 
gen flame and in the electric furnace. It is not reduced by 
hydrogen, and by carbon only with the aid of the electric 
arc. It is scarcely affected by any chemical agents at ordi- 
nary temperatures, with the exception of hydrofluoric acid 
(§ 93). When it is strongly heated with alkaline hydroxides 
or carbonates, it enters into combination with the metal, form- 
ing silicates, and these silicates are capable of dissolving silica 
at very high temperatures. When the mass cools, it constitutes 
glass, and the properties of the glass depend upon the proportions 
of silica and alkaline hydroxide or carbonate employed. If there 
be a large proportion of the alkali, the glass is soluble in water, 
and soluble glass is made by fusing silica with either potassium or 
sodium carbonate, — generally the latter, because it is cheaper. The 
solution of this substance hardens as the water evaporates, and is 
employed as a cement and in making artificial stone. 

218. Ordinary glass is made by melting in large clay pots, or 
in furnaces of peculiar construction, a mixture of fine white sand, 
sodium carbonate, and lime. If it is desired that the glass shall 
not soften at a high temperature, potassium carbonate is used in- 
stead of sodium carbonate. When the bubbles of carbon dioxide, 
which are given off from the carbonate employed, have escaped 
from the pasty liquid, the workman takes out some of the molten 
metal, as it is called, on the end of a long iron tube, through which 
he blows air into this lump of soft glass ; if the lump be in a 
bottle- mould, the glass takes the form of the mould. Common 
window-glass is made by blowing large globes which are drawn 
out into cylinders by their own weight as they hang on the blow- 
pipe. The cylinders when cold are cut open their whole length, 
and are then heated in a furnace, and when soft enough are unrolled 
into sheets. 

219. Crystal, the very heavy and perfectly colorless glass 
from which cut-glass objects are made, contains lead silicate, which 
is formed by adding red lead to the mixture of alkaline carbonate 
and sand before fusing it. A little lead is often used in making 
common glass. The dark-green color of bottle-glass is caused by 



142 LESSONS IN CHEMISTRY. 

the presence of iron in the sand used, and, in general, colored 
glasses owe the color to the presence of certain metals, as we shall 
in time learn. Plate-glass is cast on polished metallic tables, and 
while still soft is rolled out by heavy rollers, as dough is rolled. 
It is afterwards ground flat and polished by machinery. Tumblers, 
goblets, and like objects are made by pressing the soft glass into 
moulds. 

220. By the action of energetic acids the alkaline metal is at 
once removed from soluble glass. We pour into a saucer some 
thick solution of sodium silicate (sodium soluble glass), and on 
the surface of this we carefully pour some hydrochloric acid ; on 
pouring these liquids from a little height into another saucer, as 
they run out they mix on the edge, and the silica which is sepa- 
rated from the sodium hangs on the saucer in long icicle-like 
masses. 

In a beaker glass we make a rather dilute solution of sodium 
silicate, and gradually mix it with dilute hydrochloric acid. Here 
also sodium chloride is formed by the action of the hydrochloric 
acid, but no silica is precipitated. Where is it ? It must be in 
the solution, and it exists there in the form of soluble silicic acid. 
It may be separated from the sodium chloride by a process called 
dialysis ; the sodium chloride is a crystalline body, but silicic acid 
is amorphous. When a solution containing a mixture of crys- 
talline and amorphous bodies is put in a dialyser, — which is any 
glass vessel of which the bottom is cut out and replaced by a piece 
of parchment paper firmly tied on, — and the dialyser is placed in a 
vessel of water, the crystalline substance passes through the mem- 
brane, while the amorphous body remains in the interior. Then 
whew we pour our solution containing silicic acid into a dialyser, 
and set the dialyser in a vessel of water, after a time we find the 
silicic acid alone in the water of the inner vessel. This acid prob- 
ably has the composition H 4 Si0 4 = 2H 2 + SiO 2 . If we set 
aside for a few days the beaker containing it, the whole liquid is 
converted into a jelly. The silicic acid has become an insoluble 
silicic hydrate, H 2 Si0 3 = H 2 + SiO 2 . 

221. Hydrofluosilicic Acid. — We have seen how hydrofluoric 



HYDROFLUOSILICIC ACID. 



143 



acid attacks silica (§ 93). We put into a glass flask an intimate 
mixture of calcium fluoride (fluor-spar) with fine quartz sand and 
enough sulphuric acid to make a creamy liquid. We have adapted 



tube and a delivery- 
of a tall jar where it 



to our flask a cork having a safety- 
tube which may pass to the bottom 
dips into some mercury. On 
this mercury we pour some 
water, and, as our gas must 
overcome the pressure of this 
water and the mercury, we 
pour a little mercury in the 
safety-tube (Fig. 67). We 
now gently heat our flask, 
and as each bubble of gas 
passes through the mercury 
and touches the water, a ge- 
latinous deposit of silicic hy- 
drate is produced ; we use 
the mercury in order that 
the delivery-tube may not 

become stopped by this deposit. In the reaction which is taking 
place, the hydrofluoric acid which is eliminated from the fluor- 
*par and sulphuric acid at once acts upon the silica, forming silicon 
fluoride, SiFl 4 . This is the gas which comes from the flask. 




Fig. 67. 



2CaFl* + 
Calcium fluoride. 



2H 2 SO* 



+ SiO 2 = 2CaSO± + 2H 2 -f SiFl 4 
Silicic oxide. Silicon fluoride. 



When this gas comes in contact with water, a reaction takes 

place, in which silicic hydrate, H 2 Si0 3 , is formed, together with a 

gas which dissolves in the water and is called hydrofluosilicic acid. 

It is a double fluoride of silicon and hydrogen. 

3S1F1* + 3H 2 = H 2 Si0 3 + 2fSiFl*.2HFl) 
Silicon fluoride. Hydrofluosilicic acid. 

The strong solution of this gas is a highly acid liquid, and is 
valuable as a test for the metals potassium and sodium, which it 
precipitates from solutions of their salts. To a solution of potas- 
sium nitrate we add some of our filtered liquid, and at once an 



144 LESSONS IN CHEMISTRY. 

insoluble double fluoride of potassium and silicon is precipitated, 

while nitric acid now exists in the solution. 

2KN0 3 + SiFl 4 .2HFl = 2HN0 3 + SiFl 4 .2KFl 

Hydrofluosilicic acid. Silicopotassium fluoride. 






LESSON XXVIIL 
CARBON. C = 12. 

222. When we compare together a diamond, a piece of char- 
coal, and a piece of graphite from a lead-pencil, we would not 
suppose that they have many properties in common ; much less 
would we think that they are different forms of the same sub- 
stance. Yet this is the case. They are only varieties of the ele- 
ment carbon. Charcoal is formed by the decomposition of 
various compounds of carbon, and the other two modifications 
have been obtained from solutions of carbon in molten metals. 
When charcoal is dissolved in molten iron, part of the carbon 
deposits upon cooling, usually in the form of graphite, but when 
the mass is heated to about 3000° in the electrical furnace, 
and then suddenly cooled, the carbon is partly converted into 
diamond. We may say, then, that there are three modifications 
of carbon, two crystalline (diamond and graphite), and one 
amorphous, which latter includes all the varieties of charcoal. 

223. Diamond. — This is the hardest of substances : it can be 
cut and polished only by its own dust. It is found crystallized in 
regular octahedra and forms of twelve, twenty-four, and forty- 
eight faces, and the faces are usually curved (Fig. 68). The 
most highly prized varieties are perfectly colorless, but the tints 
vary through all the shades, and some diamonds are black. Its 
density is about 3.5. It is a bad conductor of heat and elec- 
tricity, and strongly refracts light. When it. is strongly heated 
in a vacuum, it blackens, and is converted into a sort of coke. 
When strongly heated in oxygen, it burns into carbon dioxide. 

224. Graphite, or plumbago, is often called black lead. It 



CARBON. 



145 



occurs in brilliant black masses, and sometimes in six-sided plates. 
It is soft enough to be easily scratched by the finger-nail, and 
leaves a black mark on paper. Its density is 2.2, and it is a good 
conductor of heat and electricity. It burns into carbon dioxide 
when heated in air to very high temperatures. It is not usually 
perfectly pure carbon, but contains one or two per cent, of foreign 




Fig. 68. 

matters. Graphite is used in lead-pencils, and for the manufac- 
ture of crucibles : in the latter it is powdered and mixed with 
clay, which binds together the graphite. It is employed also for 
coating iron objects to prevent rusting, and as a lubricant for 
machinery. 

225. The other varieties of carbon are amorphous. 

Anthracite is a hard and brittle substance, containing from 
8 to 10 per cent, of earthy matters, and sometimes even more. 

Bituminous Coal is softer and lighter than anthracite. It 
contains from 75 to 90 per cent, of carbon, with which is com- 
bined a varying proportion of hydrogen. It is the remains of 
vegetable substances which were buried in the earth in the early 
geological ages. When it is strongly heated out of contact with 
the air, various compounds of hydrogen and carbon are formed, 
together with some water and ammonia. Certain of these com- 
pounds of carbon and hydrogen are gases, and, since they contain 
only combustible elements, they are themselves combustible. 

We introduce some fragments of bituminous coal into a small 
glass retort, to the beak of which we have adapted a little jet 
(Fig. 69). When we heat the retort by a flame, heavy vapors are 
disengaged from the coal ; some of them condense in the neck of 

10 



146 



LESSONS IN CHEMISTRY. 



the retort, but those which are gaseous at ordinary temperatures 
pass out at the jet, and when we apply a flame they burn with a 
bright light. This is precisely the operation which is conducted 

for the manufacture of illuminating 
gas from bituminous coal. The coal 
is heated in clay or iron retorts (Fig. 
70), and the liquid products are con- 
densed by passing through a cold pipe ; 
since coal always contains sulphur and 
nitrogen, some hydrogen sulphide and 
ammonia are formed, and these must be 
separated from the gas, for their com- 
bustion would render the air of a room quite unwholesome. They 
are removed by passing the gas through a tall, upright pipe in 




Fig. 69. 




A, retorts. B, hj'draulic main for 
condensation of liquid products. 
C, scrubbers, in which gas is washed with 
water-spray. D, lime-purifier. E, gas- 
holder. 



Fig. 70. 



which little jets of water are playing, and the ammonia and hy- 
drogen sulphide are in great part dissolved ; the gas is still further 
purified from sulphur by passing through slaked lime, and it is then 
conducted into large gas-holders, from which it passes into the pipes 



CARBON. 147 

for consumption. The liquid which first condenses from the gas 
separates into two layers ; one is an impure solution of ammonia, 
and, together with the water used for washing the gas, forms the 
source of the ammonia of commerce ; the other is tarry, and con- 
tains numerous liquid and solid compounds of carbon and hydrogen, 
which we must study at another time. The black substance which 
remains in the clay retorts, as it does in our glass retort, is coke. 
As some of the compounds of hydrogen and carbon which are 
formed are decomposed by the high temperature of the retorts, 
these vessels become lined with the carbon, which separates, and 
forms a dense, hard, strong, and sonorous layer. It is called gas 
carbon, and is used for the carbon plates of voltaic batteries and 
for the carbon electrodes in the electric light. 

The minerals lignite and jet, of which ornaments are made, are 
varieties of bituminous coal. 

226. Charcoal is derived both from wood and from animal 
matters. Wood-charcoal was formerly made by closely piling the 
wood and covering the pile with earth, some holes being left for 
the admission of air. The combustion of part of the wood then 
produces sufficient heat to convert the rest into charcoal, or car- 
bonize it. This process is very wasteful : not only is a large quan- 
tity of wood unnecessarily burned, but the many other products, 
tar, acetic acid, and wood alcohol, which might be obtained, are 
lost. Another process is now being everywhere adopted, in which 
the wood is heated in iron retorts, and the vapors given off are 
passed through pipes surrounded by other pipes through which 
flows a stream of cold water ; the liquid products are thus con- 
densed, and the gases are conducted under the retort (Fig. 71). 
The gas from one retort is sufficient to carbonize the wood in 
another, so that after starting the operation little or no fuel is 
required and nothing is lost. 

Charcoal is brittle and sonorous; its density is about 1.5. It 
is a poor conductor of heat and electricity. It is not pure carbon : 
its combustion leaves one or two per cent, of earthy matter, prin- 
cipally the carbonates of potassium and calcium. 

Animal charcoal is made by strongly heating waste horn, 



148 



LESSONS IN CHEMISTRY. 



bone, blood, hide, and other animal matters, in closed vessels. 
When made from bone, it is called bone-black or ivory -black : it 




Fig. 71. 

then naturally contains the mineral matters of the bones, calcium 
phosphate and carbonate. These may be dissolved out by wash- 
ing the bone-black, first 
with hydrochloric acid, 
and afterwards with 
water ; it is then called 
purified animal char- 
coal. 

Lamp-black, so 
much used for the manu- 
facture of printing-ink, 
india-ink, and black 
paint, is made by burn- 
ing oil, turpentine, or 
rosin in an insufficient 
supply of air. The 
operation is conducted 
in a small furnace of 
which the chimney 
opens into a room on whose walls the thick smoke of lamp-black 
settles. Generally these walls are hung with canvas, from which 




Fig. 72. 



CARBON. 



149 



the lamp-black is removed by a conical scraper which can be 
lowered by a rope passing over a pulley (Fig. 72). Lamp-black- 
is a fine powder, and usually contains oily and tarry matters from 
the rosin : it may be purified by heating it red hot in a covered 
crucible. 

227. Properties of Charcoal. — In addition to the peculiari- 
ties of each variety which we have considered, all of the forms of 
charcoal are exceedingly porous, and they are able to absorb many 
times their volume of certain gases. We fill a rather wide glass 
tube with mercury, and, after inverting it in 
a vessel of mercury, we pass into it some am- 
monia gas, made by boiling a little ammonia- 
water in a flask. Now we heat a piece of 
charcoal red hot, to drive out all of the gases 
which it has absorbed from the air, and we 
push it under the mercury into the tube 
(Fig. 73). It rises to the surface, and in- 
stantly we see the volume of ammonia dimin- 
ishing; the gas is being absorbed by the 
charcoal, and after we remove the latter from 
the tube we will find it much heavier, and 
having a strong odor of ammonia. Charcoal 
will absorb about ninety times its volume of ammonia, fifty-five of 
hydrogen sulphide, and large quantities of most other gases. It 
absorbs less than twice its volume of hydrogen, and about eight 
times its volume of either carbon monoxide, oxygen, or nitrogen. 
These gases are driven out when the charcoal is heated. We can 
now understand that in some cases charcoal is an excellent disin- 
fectant : if a dead mouse or other small animal be buried in a box 
of powdered charcoal, it will be found after some weeks to have 
dried up, and the charcoal will have absorbed all of the unpleasant 
odors. Charcoal has also the property of absorbing many color- 
ing matters in its pores, where they are probably oxidized : animal 
charcoal possesses this property in the most marked degree. We 
pour some litmus solution into a filter containing animal charcoal, 
and the liquid passes through colorless. This peculiarity renders 




150 LESSONS IN CHEMISTRY. 

animal charcoal valuable for decolorizing many liquids, and enor- 
mous quantities of it are employed for decolorizing sugar. The 
brown color is removed from crude sugar by dissolving it in water 
and filtering the syrup through animal charcoal. An excellent 
filter for the purification of drinking-water consists of a layer of 
charcoal between two layers of sand : the charcoal must be changed 
from time to time, or it may be removed and heated red hot in a 
covered vessel ; it is then again fit for use. 

228. The strongest afiinity of carbon is for oxygen, but this 
affinity is not manifested at ordinary temperatures. When, how- 
ever, the temperature is raised to redness and the combustion of 
charcoal begins, sufficient energy is developed by the chemical 
action to keep the temperature at the combining point, and the 
oxidation goes on without further aid. The product of the com- 
bustion of carbon in air or in oxygen is carbon dioxide, CO 2 . Be- 
cause of its strong affinity for oxygen, charcoal can remove that 
element from various oxidized bodies : it is a reducing agent. We 
have already seen an example of this reduction when we heated 
charcoal with cupric oxide (§ 13). When such a reduction re- 
quires a temperature about redness, carbon dioxide is formed, and 
this was the case with cupric oxide and charcoal. 

2CuO + C = Cu2 + CO 2 
Cupric oxide. Copper. Carbon dioxide. 

When, however, the reduction requires a white heat, or near 
that temperature, carbon monoxide, CO, is formed. This is the 
action of charcoal on zinc oxide. 

Zn 



ZnO + C 

Zinc oxide. 




CO 

Carbon monoxide. 




LESSON 
OXIDES OF 


XXIX. 
CARBON. 



229. Carbon Monoxide, CO. — This gas might be made by 
heating to whiteness in clay retorts a mixture of zinc oxide and 
charcoal, but this would require a furnace and be inconvenient. 



CARBON MONOXIDE. 



151 



We can prepare the gas more conveniently by heating in a glass 

flask a mixture of oxalic and strong sulphuric acids. The oxalic 

acid, which is a compound of carbon, oxygen, and hydrogen, is 

then decomposed into carbon monoxide, carbon dioxide, and water. 

CWH* = CO 2 + CO + H 2 

Oxalic acid. Carbon dioxide. Carbon monoxide. 

The water will be retained by the sulphuric acid, but we must pass 
the gases through a bottle containing a solution of sodium hydrox- 
ide, by which the carbon dioxide will be absorbed. Sodium car- 
bonate is formed in the bottle, and we collect the carbon monoxide 
over water (Fig. 74). 




Fig. 74. 



230. Carbon monoxide is a colorless, odorless gas. Its density 
compared to air is 0.967, or compared to hydrogen, 14. It is in- 
soluble in water. It is very poisonous : when it is taken into the 
lungs it combines with the red globules in the blood, and prevents 
them from carrying into the system the oxygen which is necessary 
for the processes of life (§ 33). 



152 LESSONS IN CHEMISTRY. 

Just as we made a similar experiment with hydrogen, carefully 
we lift our jar from the water, and push into it a lighted taper. 
The gas takes fire and burns with a blue flame at the mouth of 
the jar, but the taper is extinguished. Carbon monoxide will not 
support combustion, but it will combine with oxygen to form carbon 
dioxide. 

It is interesting here to study the amount of heat disengaged by the com- 
bustion of carbon. Naturally, we can understand that when a given weight 
of charcoal is burned, a fixed quantity of heat will be developed, enough to 
raise a certain weight of water, let us say, from 0° to 1°. When this same 
weight of charcoal is converted into carbon monoxide, the combustion of the 
latter gas will not heat as much water through the same temperature. What 
has become of the energy which has disappeared from the charcoal ? It must 
have been lost from the atom of carbon and the first atom of oxygen with 
which it combined. Many careful experiments have shown that this is the 
case, and that the same quantity of carbon always develops the same quan- 
tity of heat when it is converted into carbon dioxide, whether it is so con - 
verted at once, or whether it first forms carbon monoxide and this combines 
with an additional atom of oxygen. We so have a method which enables us 
to determine the heat or energy of formation of bodies like carbon monoxide. 
For if from the amount of energy developed by the conversion of a certain 
amount of carbon into carbon dioxide, we subtract that which is developed by 
the combustion of a quantity of carbon monoxide containing the same weight 
of carbon, we will have the energy with which one atom of carbon combines 
with one atom of oxygen. In making such determinations, a number of 
grammes of the substance is taken which would express the atomic weight if 
one atom of hydrogen weighed one gramme ; and the result, which is expressed 
in the number of kilogrammes of water which would be raised from 0° to 1° 
by the heat produced, is called the heat of formation of the compound. 

231. Carbon monoxide is formed by the action of carbon dioxide 
on carbon at very high temperatures. 

CO 2 + C = 2CO 
When fresh coal is thrown on a hot fire, the escape of the carbon 
dioxide from the burning coal is retarded, and that gas remains 
in contact with the coal long enough to be partially reduced to 
carbon monoxide. The latter then occasions the blue flame with 
which we are all familiar. Carbon monoxide has the property of 
passing through the pores of red-hot iron, and it often so escapes 
through the iron of stoves which are not properly lined with fire- 
brick ; fortunately, the gas formed from coal has a decided odor, 



CARBON DIOXIDE. 



153 



usually due to the sulphur in the coal, and this odor generally 
makes us aware of the presence of the poisonous gas. 

232. When steam is passed over hot coal or charcoal, a mixture 
of carbon monoxide and hydrogen is formed. 
C + H20 = CO + H2 
When this mixture is passed through volatile compounds of 
hydrogen and carbon, such as we shall learn are contained in the 
lighter kinds of petroleum, the gases become charged with the 
vapors of those compounds : if they then be passed through hot 
pipes, various gaseous compounds of carbon and hydrogen are 
formed, and these burn with light-giving flames. The water-gas 
used for illumination in some cities is manufactured in this manner. 

233. Carbonyl, Chloride. — Carbon monoxide combines directly with chlo- 
rine when a mixture of the two gases is exposed to direct sunlight, and a 
suffocating gas, phosgene or carbonyl chloride, COG 2 , is formed. The carbon 
monoxide acts as a radical, which is called carbonyl. Carbonyl chloride reacts 
with water, and yields carbon dioxide and hydrochloric acid. 

COC12 + H 20 = CO* + 2HC1 

Carbon monoxide also unites directly with certain metals. When it is passed 
over finely divided nickel, there results nickel carbonyl, Ni(CO)*, a volatile, 
highly-refracting liquid. Its vapor is very poisonous, and it burns with a 
luminous flame. A similar iron compound, Fe(CO) 5 , is also known. 

234. Carbon Dioxide, CO 2 . — In a gas-bottle, like that which 
we used for preparing hydrogen (Fig. 75), we put some water and 

broken marble, and pour hydro- 
chloric acid through the funnel- 
tube. As the gas with which 
we wish to experiment is very 
heavy, we collect it by down- 
ward dry displacement, passing 
the end of our delivery-tube to 
the bottom of the jar. When 
the effervescence in the bottle 
has continued for a moment, 
we put a lighted taper into the 

iar in which we are collecting 
Fig. 75. J & 

the gas : the name is at once 

extinguished. The jar is full of carbon dioxide, and for such 




154 LESSONS IN CHEMISTRY. 

substances as the matter of the taper the oxygen in this gas 

has exhausted its energy in combining with the carbon : it will 

unite with no more. In our gas-bottle we have formed water and 

calcium chloride, for marble is calcium carbonate. 

CaCO 3 + 2HC1 = CaCl 2 + H 2 + CO 2 
Calcium carbonate. Calcium chloride. 

235. Carbon dioxide is a colorless gas ; it has a faint but some- 
what pungent odor and taste. Its density compared to air is 
1.529, or compared to hydrogen, *22. We balance on a scale-pan 
an open and erect paper bag, into which we quickly pour the car- 
bon dioxide from our jar : the descending pan at once shows us 
that the gas is heavier than the air which it displaces. At 0°, it 
is converted into a colorless liquid by a pressure of thirty-six at- 
mospheres, and when the liquid is allowed to evaporate rapidly, 
it absorbs so much heat in assuming the gaseous state that the 
temperature falls to — 78°, and a part of the carbon dioxide is 
frozen to a snow-like mass. When touched, this solid produces a 
burn like fire, for the life of animal tissues cannot continue at 
such low temperatures. 

Carbon dioxide is soluble in its own volume of water, and the 
quantity of the gas which can be absorbed by a given quantity of 
water is directly proportional to the pressure : if the pressure be 
doubled, twice as much gas will be dissolved, etc. We know, 
however, that by a double pressure two volumes of any gas will 
be reduced to one volume (Mariotte's law) : hence we may say 
that water always dissolves its own volume of carbon dioxide, no 
matter what the pressure. When the pressure is diminished or 
removed, the gas escapes with effervescence, until the volume 
remaining dissolved is equal to that of the water. The beverage 
generally known as aerated water or soda-water, is simply water 
into which about five times its volume of carbon dioxide has been 
pumped. 

236. We have already seen that carbon dioxide neither burns 
nor supports combustion, and a simple experiment shows us its 
power of extinguishing burning bodies, at the same time that we 
will be reminded of its weight. We fix a short taper or piece of 
candle in a cork, and, after lighting it, put it in a small jar into 



CARBON DIOXIDE. 



155 



which we pour the carbon dioxide from another jar which we have 
filled by dry displacement : the flame is extinguished as if we had 
poured water on it (Fig. 76). Carbon 
dioxide is not poisonous, but it pro- 
duces death by suffocation, — that is, 
exclusion of oxygen. It collects in 
wells and brewers' vats, and may then 
be detected by its power of extinguish- 
ing flames lowered into it : if there be 
enough of the gas present to extin- 
guish or nearly extinguish a flame, 
it would not be safe for a man to enter 
such a place before removing the gas, 
which can be done by agitating the 
air so that currents maybe established. 

237. Certain substances which have the power of reducing 
carbon dioxide may burn in the gas. Over a piece of the metal 
potassium contained in a glass bulb, we pass carbon dioxide that 





Jmg. 



has been dried by passing through a tube containing pumice-stone 
and sulphuric acid. When we warm the potassium, it takes fire 
and burns with a red light, and a deposit of carbon is formed 
in the bulb (Fig. 77). The potassium has reduced some of the 



156 LESSONS IN CHEMISTRY. 

carbon dioxide, and has combined with another portion, forming 
potassium carbonate. 

238. We pass a few bubbles of carbon dioxide into some lime- 
water : the liquid quickly becomes milky by the formation of in- 
soluble calcium carbonate. This test enables us to recognize carbon 
dioxide. Calcium carbonate is insoluble in water, but if we pass 
the gas for a long time through our milky liquid, the cloudiness 
disappears ; water containing carbon dioxide in solution will dis- 
solve calcium carbonate. If, however, we boil this liquid so that 
all of the dissolved carbon dioxide shall be driven out, the calcium 
carbonate again separates, usually as small crystalline particles, 
which settle as the water cools. The stalactites and incrustations 
in caves are formed by the drippings of water holding calcium 
carbonate in solution by an excess of carbon dioxide ; as the gas 
passes off gradually into the air, the calcium carbonate becomes 
insoluble and is deposited. 



LESSON XXX. 

CARBONATES. 

239. Carbon dioxide corresponds to an acid which would be formed by the 
action of one molecule of the gas on one molecule of water. The solution of 
the gas in water is feebly acid, and we may believe that it contains carbonic 
acid, H 2 C0 3 = CO 2 f H 2 0. Carbon dioxide is often called carbonic anhydride ; 
anhydride means without water, and carbonic anhydride thus signifies carbonic 
acid less the elements of water. Although this acid cannot be separated from 
the solution, for when the water is evaporated carbon dioxide is driven out, 
there are numerous salts formed by the replacement of one or both of the 
atoms of hydrogen in H 2 C0 3 . These salts are the carbonates; they maybe 
easily recognized by the action of hydrochloric acid, which produces with 
them an effervescence due to the escape of carbon dioxide; we may identify 
this gas by the milkiness which it produces in lime-water. 

There are two classes of carbonates ; those in which both atoms of hydrogen 
in H 2 C0 3 are replaced by metal, and others in which only one of these atoms 
is replaced. The latter are called the acid carbonates, or sometimes the dicar- 
bonates. Carbonic acid is, then, dibasic. 



SODIUM CARBONATE. 157 

With the exception of the carbonates of sodium, potassium, and 
lithium, all of the carbonates of the metals are insoluble in water : 
they dissolve slightly, however, in water containing carbon dioxide. 
They all effervesce when treated with hydrochloric or sulphuric 
acid, carbon dioxide being disengaged. 

240. Sodium Carbonate, Na 2 C0 3 . — Enormous quantities of 
this salt, which is commonly called soda, or sal soda ) are used for 
the manufacture of glass and of soap, and for the preparation of 
the many compounds of sodium that are used in the arts. It is 
usually manufactured from common salt, and the process which is 
coming into general use depends on a reaction between the salt 
and ammonium acid carbonate (§ 251). We mix saturated solu- 
tions of ammonium acid carbonate and sodium chloride, and a fine 
white deposit is formed in the liquid. This is sodium acid car- 
bonate, and ammonium chloride exists in the solution. 

NaCl + (NH*)HC03 = NH*C1 + NaHCO 3 

Sodium chloride. Ammonium acid Ammonium chloride. Sodium acid 

carbonate. carbonate. 

This operation is conducted on a large scale, and the sodium 
acid carbonate is converted into sodium carbonate by the action of 
heat. Two molecules of the acid carbonate then lose one molecule 
of water, and one of carbon dioxide. 

2XaHC0 3 = Xa 2 C0 3 + H 2 + CO 2 

Sodium acid carbonate. Sodium carbonate. 

The ammonium chloride is converted into ammonia by heating 
it with lime (§ 136), and the carbon dioxide formed by heating the 
sodium acid carbonate, together with more which is obtained from 
the gases of the furnace-chimneys, serves to convert the ammonia 
again into ammonium acid carbonate. 

NH3 + H 2 + CO 2 = (XH*)HC0 3 

The only waste product is, then, the calcium chloride left after 
the preparation of the ammonia. In the older process, which the 
ammonia-soda process is gradually replacing, the sodium chloride 
is first converted into sodium sulphate (§ 77) ; this is mixed with 
chalk and coal, and the mixture heated by the flame of a rever- 
beratory furnace (Fig. 78) yields a mixture containing calcium 
sulphide and sodium carbonate. The last is dissolved out by water, 



158 



LESSONS IN CHEMISTRY. 



and the waste heat of the furnace is employed not only to evap* 
orate the solution obtained (C and D), but to dry the mixture of 
sodium sulphate, chalk, and coal (B) before introducing it into 
the hottest part of the furnace (A). This is named, from its 
inventor, the Leblanc process. 

Sodium carbonate is also manufactured from the mineral cryolite, 
a double fluoride of sodium and aluminium, of which large quan- 




Fig. 78. 

tities are found in Greenland. The cryolite is heated with lime, 

and the reaction yields calcium fluoride and a compound known as 

aluminate of sodium : it is a combination of the oxides of sodium 

and aluminium. 

2AlE 3 .3NaF + 6CaO = 6CaF 2 + Al 2 3 .3Na 2 

Cryolite. Lime. Calcium fluoride. Aluminate of sodium. 

The aluminate of sodium is dissolved from the mass by water, and 

carbon dioxide passed through the solution forms sodium carbonate, 

while insoluble aluminium hydroxide is precipitated. 

Al 2 3 .3Na 2 + 3C0 2 + 3H 2 = 2A1(0H) 3 + 3Na 2 C0 3 

Aluminate of sodium. Aluminium hydroxide. 

The aluminium hydroxide is used for the manufacture of alum. 

Sodium carbonate forms large crystals containing ten molecules 
of water of crystallization for one molecule of the salt. When it 
is exposed to the air, it gradually loses this water, and falls to a 
dry, white powder : it is said to effloresce. The crystals dissolve 
in about four times their weight of water at 20°, and the solution 
has an alkaline reaction and an unpleasant alkaline taste^ The 
salt is insoluble in alcohol. 

241. Sodium Acid Carbonate, NaHCO 3 , is less soluble than 



POTASSIUM CARBONATE. 159 

the carbonate, and is precipitated when carbon dioxide is passed 
through a saturated solution of the latter salt. It is usually a 
white powder, which, when heated or boiled with water, is decom- 
posed into water, sodium carbonate, and carbon dioxide which 
escapes. It is commonly called bicarbonate of soda, or baking- 
soda; it forms part of the mixtures known as baking-powders, 
which contain some acid substance that may react with the acid 
carbonate, setting free carbon dioxide in bubbles through the 
dough. Sodium acid phosphate is such a substance ; sodium acid 
carbonate may convert it into either disodium or trisodium phos- 
phate. 

XaHCO 3 - XaH 2 PO± = Xa 2 HP0 4 + H 2 + CO 2 

242. Potassium Carbonate, K 2 C0 3 . — This compound is com- 
monly called potash, because it was for a long time derived only 
from wood-ashes ; it is extracted from the ashes by causing water 
to trickle through them, a process which is called lixiviation. The 
solution so obtained is evaporated to dryness, and the residue 
strongly heated in the air. The potash of commerce cou tains only 
from 60 to 80 per cent, of potassium carbonate. The remainder 
consists of other potassium salts, principally the chloride and sul- 
phate : when these are partially removed by an imperfect purifica- 
tion, the product is called pea rl-ash. 

At Stassfurt, in Prussia, there are large deposits of a double 
chloride of potassium and magnesium ; the mineral is called 
camaUite, and contains KCl.MgCl 3 -j- 6H 2 0. It is decomposed 
by boiling with water, and, on cooling, potassium chloride crys- 
tallizes from the liquid, while magnesium chloride remains in 
the solution. Potassium carbonate is now manufactured from 
this potassium chloride by a method similar to the Leblanc 
process for sodium carbonate. 

Potassium carbonate is white, and dissolves in less than its own 
weight of water. It is very alkaline, and has a burning taste. 
It is deliquescent ; that is, it attracts moisture from the air and 
dissolves in the water so absorbed. It may be obtained in crys- 
tals containing two molecules of water by allowing a hot concen- 
trated solution to cool. 



160 LESSONS IN CHEMISTRY. 

243. Potassium Acid Carbonate, KHCO 3 , is prepared, like 
bicarbonate of soda, by passing carbon dioxide through a solu- 
tion of the neutral carbonate. It is less soluble than the latter, 
and separates from the solution in crystals. Like sodium acid 
carbonate, it is decomposed by heat, whether it be dry or in 
solution. 

244. Calcium Carbonate, CaCO 3 . — This substance is one of 
the most abundant of minerals. As calotte, or Iceland spar, it 
forms doubly-refracting rhombohedra : as aragonite, it occurs 
in rhombic prisms. It also constitutes marble, limestone, chalk, 
and the greater part of the matter of shells and corals. When 
heated to bright redness in open vessels, it is decomposed into 
carbon dioxide and lime, which is calcium oxide. 

245. Strontium Carbonate, SrCO 3 , constitutes the white 
mineral strontianite. 

246. Barium Carbonate, BaCO 3 , is found crystallized in nature 
in the mineral witherite. The carbonates of calcium, strontium, 
and barium are precipitated when a solution of sodium or potas- 
sium carbonate is added to a neutral solution of any calcium, 
strontium, or barium salt. 

247. Magnesium Carbonate, MgCO 3 , constitutes the minerals 
magnesite and giobertite. Dolomite is a double carbonate of calcium 
and magnesium ; it is a magnesian limestone. White magnesia is 
a variable compound of magnesium carbonate and magnesium 
hydrate, made by adding an excess of sodium carbonate solution 
to a boiling solution of magnesium sulphate, and drying the pre- 
cipitate. 

248. Zinc Carbonate, ZnCO 3 , constitutes the mineral smith- 
sonite, an important ore of zinc. 

249. Ferrous Carbonate, FeCO 3 , is found native in brown 
crystals as siderite or spathic iron. 

250. Lead Carbonate, PbCO 3 , is met with as cerussite. It 
is precipitated as an amorphous white powder when any soluble 
lead salt is treated with sodium carbonate. White lead is a mix- 
ture of varying proportions of lead carbonate and lead hydroxide. 
It is manufactured by the joint action of carbon dioxide and 





AMMONIUM CARBONATES. 161 

vapor of acetic acid on metallic lead. Acetic acid, that is, vinegar, 
is put into earthen pots, and the lead, either in a rolled sheet or 
in flat rings, is supported on little pro- 
jections above the vinegar (Fig. 79). 
A great number of these pots are pre- 
pared, and loosely covered with disks 
of lead (D) : they are then arranged in 
layers on boards, each layer resting on 
a bed of refuse bark from tanneries, or 
of horse-manure. These substances un- 
dergo a sort of slow oxidation, and dis- 
engage carbon dioxide, which in the 

presence of the acetic acid converts the surface of the lead into 
carbonate. We may suppose that lead acetate is first formed, and 
that this is at once decomposed by the carbon dioxide, forming 
lead carbonate, while acetic acid is regenerated. When the greater 
part of the lead has been so changed into carbonate, the pots are 
opened and the white lead is scraped from the remaining metal. 

There is another process, which depends on the facility with 
which lead acetate solution dissolves lead oxide, and with which 
this oxide is precipitated as carbonate when carbon dioxide is passed 
through the solution. Lead acetate is then again formed in the 
solution, and is used to dissolve more oxide, which is in its turn 
precipitated by carbon dioxide. 

Excepting sodium carbonate and potassium carbonate, all the 
other salts which we have just considered are decomposed by heat 
into carbon dioxide and oxide of the metal. 

251. Ammonium Carbonates. — The ammonium carbonate of 
commerce is commonly called a sesquicarbonate, and is probably 
a mixture of several bodies. It is made by subliming a mixture 
of chalk and ammonium sulphate in large retorts. Its compo- 
sition is expressed by the formula 2[(NH*) 2 C0 3 ] + CO 2 + H 2 0. 
It is a crystalline substance, having a strong ammoniacal odor 
and a sharp, burning taste. It is soluble in water. When it is 
exposed to the air, it gradually decomposes, losing ammonia and 
leaving ammonium acid carbonate, NH 4 .HC0 3 . The latter body 

n 



162 LESSONS IN CHEMISTRY. 

may also be formed by passing carbon dioxide into ammonia-water 
until no more of the gas is absorbed. It then crystallizes when 
the liquid is cooled. Ammonium carbonate, (NH 4 ) 2 C0 3 , separates 
in crystals when the ammonium carbonate of commerce is dis- 
solved in ammonia-water and the solution is artificially cooled. 



LESSON XXXI. 

COMPOUNDS OF CARBON WITH SULPHUR AND 
NITROGEN. 

252. Carbon Disulphide, CS 2 . — When sulphur vapor is passeQ 
over red hot charcoal or coal, the two elements combine, forming 
carbon disulphide, which passes off as a vapor that may be con- 
densed in any suitable cooling apparatus. This substance is 
manufactured by throwing sulphur into coal heated to redness in 
inclined iron cylinders, provided with openings for the introduction 
of sulphur and the escape of the vapor. 

It is a colorless liquid, having the property of highly refracting 
light. When pure, it has a rather pleasant odor, but the com- 
mercial product usually contains small quantities of other com- 
pounds which communicate to the liquid a strong and often dis- 
gusting odor : it is purified by distillation with lime. Its density 
at 15° is 1.27. It is almost insoluble in water. It boils at 46°. 
It is very inflammable : we heat a wire to redness, withdraw it 
from the flame, and for some time after it has cooled below a 
visible heat it will still inflame carbon disulphide contained in a 
small dish. The vapor forms an explosive mixture with the air 
or with oxygen. The products of the combustion of carbon di- 
sulphide are carbon dioxide and sulphur dioxide. 
CS 2 + 30 2 = CO 2 + 2S0 2 

If a few thin iron wires are held in the flame of vapor of carbon 
disulphide which is heated in a small test-tube provided with a 
cork and jet, the air and sulphur combine, producing brilliant 



CARBON BISULPHIDE. 



163 




Fig. 80. 



sparks and molten globules of iron sulphide, which drop from the 
ends of the wires (Fig. 80). 

Carbon disulphide is used 
for extracting oil from seeds 
and other matters, for it 
dissolves fatty substances 
quickly and in the cold, 
and may readily be distilled 
and recovered from the so- 
lution, leaving a much larger 
quantity of oil than could 
be extracted by pressure. 
It is used in vulcanizing 
caoutchouc, an operation 
which depends on the com- 
bination of the caoutchouc 
with a certain quantity of 

sulphur, as it is able to dissolve not only the caoutchouc but the 
sulphur chloride which is used in the operation. We have seen 
that carbon disulphide dissolves iodine, sulphur, and phosphorus. 

253. Carbon disulphide is closely related to carbon dioxide. 
We have studied the general composition of the carbonates, and 
know that they correspond to a carbonic acid which should con- 
tain H 2 C0 3 . There is also a series of sulpho-carbonates, exactly 
similar to the carbonates in composition, but they contain three 
sulphur atoms instead of three atoms of oxygen. 

(H 2 C0 3 ), Carbonic acid. H 2 CS 3 , Sulphocarbonic acid. 

Na 2 C0 3 , Sodium carbonate. Na 2 CS 3 , Sodium sulpho-carbonate. 

K 2 C0 3 . Potassium carbonate. K 2 CS 3 , Potassium sulpho-carbonate. 

The sulpho-carbonates as well as carbon disulphide are em- 
ployed to destroy vermin, a purpose for which they are quite 
effective. 

254. If we compare together the molecules of the few compounds of carbon 
which we have studied, we will find that in all excepting one an atom of car- 
bon has as much combining power as four atoms of hydrogen. It is tetra- 
tomic : in carbon dioxide, because it is united with two atoms of oxygen, each 
of which is worth two atoms of hydrogen j in sodium carbonate, for it is there 
combined with one atom of oxygen and two other atoms of oxygen, each of which 
brings into the system a sodium atom as a satellite; in carbon disulphide, 



164 



LESSONS IN CHEMISTRY. 



where it is combined with two atoms of diatomic sulphur; and there is also 
a carbon oxy sulphide, COS, in which one atom of oxygen and one of diatomic 
sulphur satisfy the combining capacity of the tetratomic carbon atom. How- 
ever, in carbon monoxide either the carbon atom must be diatomic, — that is, 
worth two atoms of hydrogen, — or the oxygen atom must be tetratomic. Since 
we know that carbon monoxide can combine with two atoms of chlorine, each 
of which has the power of one hydrogen atom, and that it may also combine 
with another oxygen atom, we must consider that the carbon atom is diatomic 
in a molecule of carbon monoxide, which is then an unsaturated compound. 
We may represent the structure of these molecules, as we have expressed that 
of others, by structural formulae : 



c=o 


Cl-C-Cl 


0=C=0 


NaO-C-ONa 


Carbon 
monoxide. 




Carbonyl 
chloride. 


Carbon 
dioxide. 




Sodium 
carbonate. 



S=C=S 



0=C=S 



Carbon Carbon 

disulphide. oxysulphide. 

After a time we shall become acquainted with compounds in which the 
group CS acts as a radical, precisely as does the group carbonyl, CO ; so that 
we may consider the compound COS either as a combination of carbonyl with 
an atom of sulphur, or as a compound of the group CS with one atom of 
oxygen. 

255. In the combining power of its atoms, silicon resembles carbon. It also 
is tetratomic, as we can understand from the composition of the silicon com- 
pounds which we have studied ; but there is no monoxide corresponding to 
carbon monoxide, and silicon does not appear to be diatomic in any com- 
pounds. 

F-Si-F 



0=Si=0 



HO-Si-OH 

II 




A 
FF 



Silicic oxide. Silicic hydrate. Silicon fluoride. 

256. Cyanogen, C 2 N 2 . — We put into a test-tube some mer- 
curic cyanide, a white and 
very poisonous compound of 
mercury, carbon, and nitrogen; 
then we adapt to our test-tube 
a cork in which we have fitted 
a bent tube bearing a little 
bulb containing some small 
pieces of the metal potassium 
(Fig. 81); the outer end of 
this tube is drawn into a fine 
jet. We now heat the mer- 
curic cyanide, and presently metallic mercury begins to deposit in 




Fig. 81. 



CYANOGEN. 165 

the cooler part of the tube, and a gas is escaping from the jet : 
we light it, and it burns with a beautiful peach-blossom-colored 
flame. 

This gas is cyanogen. It is a colorless gas, having an odor 
like that of bitter almonds, and is quite poisonous. Its density 
compared to air is 1.8064, or compared to hydrogen, 26; its 
molecular weight is, then, 52, and analysis has shown that it con- 
tains carbon and nitrogen in the proportions indicated for one 
atom of each. Since its molecular weight is 52, a molecule of 
cyanogen gas must contain two atoms of carbon and two of nitro- 
gen. Cyanogen is converted into a liquid by pressure or by a 
temperature of — 20.7°. It dissolves in about one-quarter its 
volume of water, but the solution soon decomposes, and then 
always contains ammonia or some ammonium compound. The 
combustion of cyanogen produces nitrogen and carbon dioxide. 

257. By the aid of a spirit-lamp, we now heat the bulb con- 
taining the potassium. There is a bright flash of light : the 
potassium and cyanogen have combined, and formed potassium 
cyanide. Cyanogen is, then, capable of entering into combination. 
When, however, we analyze the potassium cyanide formed, we 
find that it contains potassium, carbon, and nitrogen in the pro- 
portions required for one atom of each ; its formula is KCN. 
The cyanogen molecule C 2 N 2 has then separated into two groups 
CN, each of which has combined with an atom of potassium. 
The group CN is a radical, and free cyanogen resembles free chlo- 
rine in this respect, for the molecule of chlorine contains two 
atoms, while the molecule of cyanogen contains two groups or 
radicals. 

Cl-Cl NC-CN 

The reaction between potassium and cyanogen is, then, like that 
between potassium and chlorine ; both are double decompositions. 

K-K + Cl-Cl = KCl + KC1 

K-K + (CN)-(CN) = KCN + KCN 
258. In free cyanogen gas is it the carbon atoms which are united together, 
or the nitrogen atoms ? When cyanogen or its compounds decompose, the ni- 
trogen atoms always form compounds in which they are triatomic, having the 
combining power of three atoms of hydrogen. Then we must believe that they 



166 



LESSONS IN CHEMISTRY. 



are also triatoniic in cyanogen, and since the carbon atom is worth four hydro- 
gen atoms and the nitrogen atom only satisfies three-fourths of this combining 
power, the carbon atoms must be united together, and we consequently write 
cyanogen gas NiLC-CEN or (CN) 2 . We see also that the potassium atom in 
potassium cyanide, and the mercury atom in mercuric cyanide, must be united 
to the carbon atom of cyanogen. (Compare §§ 262 and 334.) 
K-CIN NlC-Hg-CEN 

Although carbon and nitrogen do not combine directly, potassium cyanide 
is formed when either nitrogen gas or ammonia is passed over a highly-heated 
mixture of charcoal with potassium carbonate or potassium hydrate. All of 
the compounds of cyanogen are prepared from potassium ferrocyanide ($ 266). 



LESSON XXXII. 



HYDROCYANIC ACID.— CYANIDES. 



259. Hydrocyanic Acid, HCN. — This dangerous poison, com- 
monly called prussic acid, is formed when a cyanide is treated 
with a dilute acid ; as by the action of hydrochloric acid on mer- 
curic cyanide. 

Hg(CN) 2 + 2HC1 = HgCl 2 + 2HCN 
Mercuric cyanide. Mercuric chloride. 

It is usually made by distilling 8 parts of potassium ferrocya- 
nide with a cooled mixture of 9 parts of sulphuric acid and 14 
parts of water. The beak of the retort containing this mixture is 

inclined upwards, in order 
that the water may condense 
and run back into the retort ; 
the vapor of hydrocyanic acid 
is dried by passing through a 
calcium chloride tube placed 
in water heated to about 30°, 
and then condensed in a flask 
surrounded by a mixture of ice and salt (Fig. 82). 

260. Hydrocyanic acid is a colorless, very volatile liquid ; its 
odor resembles that of bitter almonds. Its density is about 0.7 ; 




Fig. 82. 



HYDROCYANIC ACID. 167 

it freezes at — 15°, and boils at 26.5°. The density of its vapor 
compared to hydrogen is 13.5, corresponding exactly with the 
molecular weight implied by the formula HCN. It dissolves in 
all proportions of water, and a two per cent, solution is used in 
medicine. It is combustible, and when ignited burns into water, 
carbon dioxide, and nitrogen. It is exceedingly poisonous, and 
the accidental inhalation of its vapor has in some cases proved 
fatal. 

261. It is often important to be able to recognize hydrocyanic 
acid ; and we may do so by the following tests. We make our 
solution of hydrocyanic acid for these tests by adding a little dilute 
sulphuric acid to some solution of potassium cyanide. The liquid 
then contains hydrocyanic acid and potassium sulphate. 

Over the beaker glass in which we have prepared this solution, 
we invert a watch-glass or glass plate on which we 
have placed a drop of silver nitrate solution (Fig. 
83) : this drop soon becomes clouded from the for- 
mation of insoluble silver cyanide ; the white de- 
posit does not darken quickly on exposure to light, and 
leaves metallic silver when heated ; these characters 
distinguish it from silver chloride, which would be yiq. 83. 
formed if the liquid contained hydrochloric acid. 

We now invert over our beaker another watch-glass containing 
a drop of ammonium sulphide which has become yellow by expo- 
sure to light and air : some ammonia has escaped from it, and it 
contains an excess of sulphur. In a little while this drop becomes 
colorless : a compound called ammonium sulphocyanate has been 
formed in it. 




(NH*) 2 S 
immouium 


+ 


S 2 


+ 2HCN - 


= 2NH*SCN + 
Ammonium 


H 2 S 


sulphide. 








sulphocyanate. 





If we now carefully warm the spot until it no longer has the odor 
of hydrogen sulphide, and then touch it with a drop of ferric 
chloride solution, a blood-red color appears. This color is due to 
the formation of ferric sulphocyanate (§ 277). 

We mix in a test-tube a few drops of our hydrocyanic acid 



168 LESSONS IN CHEMISTRY. 

solution with a little ferrous sulphate and ferric sulphate, and add 
a little strong solution of sodium hydroxide : a dirty deposit forms, 
but when we add an excess of hydrochloric acid, a part of the 
deposit is dissolved, and a fine blue precipitate, Prussian blue 
(§ 267), remains. 

262. Hydrocyanic acid does not keep long, soon decomposing, whether it be 
pure or in solution. It undergoes an interesting reaction with strong hydro- 
chloric acid, and the reaction is more interesting because it is characteristic 
of all the cyanides. When we mix hydrocyanic acid with strong hydro- 
chloric acid, the mixture becomes hot, and a mass of crystals of ammonium 
chloride separate. The most curious part of this reaction is, that it takes 
place between the hydrocyanic acid and the water of the hydrochloric acid; 
the nitrogen atom of the former is exchanged for an atom of oxygen and a 
hydroxyl group. 

HCN + 2H20 = HCO.OH + NH3 

The ammonia formed combines with the hydrochloric acid. The compound 
HCO.OH is called formic acid. When a solution of potassium cyanide is 
boiled, it is converted into potassium formate by a similar reaction. 

KCN + 2H20 = HCO.OK + NH^ 

Potassium cyanide. Potassium formate. 

All the acids of carbon which we shall presently have occasion to study, 
may be formed by the replacement of the nitrogen atoms of corresponding 
cyanides by an oxygen atom and a hydroxyl group. 

263. Potassium Cyanide, KCN, is made by heating dry po- 
tassium ferrocyanide red hot in earthen retorts. After the mass 
cools, it is extracted with alcohol, and when the filtered liquid is 
evaporated it leaves a white mass of potassium cyanide. It is 
very soluble in water, and may be crystallized in cubes. Upon 
heating with- sulphur, it is converted into potassium sulpho- 
cyanate, KS.CN. Solutions of potassium cyanide dissolve the 
cyanides of gold, silver, zinc, and other metals, forming double 
cyanides, and some of these also result when the free metals, 
in a finely divided state, are treated with potassium cyanide. 
Hence this salt is extensively used in the extraction of the 
precious metals from their ores, in photography, and in electro- 
plating. Potassium cyanide is exceedingly poisonous. 

264. Silver Cyanide, AgCN, is formed as a white precipitate 
when a solution of silver nitrate is treated with the exact quan- 
tity of potassium cyanide required for one molecule of each. 
When heated, it decomposes into silver and cyanogen gas. 



CYANIDES. 169 

265. Mercuric Cyanide, Hg(CN) 2 . — This compound maybe 
made by dissolving mercuric oxide in dilute hydrocyanic acid, 
but it is usually prepared by boiling a mixture of one part of 
potassium ferrocyanide, two parts of mercuric sulphate, and 
eight parts of water. The mixture is filtered while boiling, and 
mercuric cyanide separates from the filtrate in colorless, anhy- 
drous, square prisms. It dissolves in eight times its weight of 
cold water. 

266. Potassium Ferrocyanide, K 4 Fe(CN) 6 . — Potassium fer- 
rocyanide is the starting-point for the preparation of other com- 
pounds of cyanogen. There are a number of processes for its 
manufacture : the most common of them consists in heating waste 
animal matters containing nitrogen, such as blood, horn, scraps of 
skin and leather, with potassium carbonate and scrap iron. After 
the mass has cooled, it is exhausted with boiling water, and the 
concentrated solution deposits the ferrocyanide in crystals. 

These crystals are yellow, and contain three molecules of water 
of crystallization, which may be driven out by a temperature of 
100° ; the anhydrous salt then remains as a white powder. Crys- 
talline potassium ferrocyanide, which is commonly called yellow 
prussiate of potash, dissolves in four times its weight of cold or 
twice its weight of boiling water, and is insoluble in alcohol. It 
is not poisonous. 

The group of atoms Fe(CN) 6 which it contains is a radical, and 
takes part in double decompositions without undergoing change. 
There is a hydroferrocyanic acid, H 4 Fe(CN) 6 . We add some 
cupric sulphate to solution of potassium ferrocyanide, and a ma- 
hogany-colored precipitate of cupric ferrocyanide is formed, while 
potassium sulphate goes into solution. 

2CuS0 4 + K*Fe(CN)6 = 2R2S0 4 + Cu2Fe(CN) 6 

Solution of potassium ferrocyanide causes the formation of 
insoluble ferrocyanides in solutions of many metallic salts, and 
the color of the precipitate is a means frequently employed for 
identifying the metals. With zinc sulphate, we would have zinc 
ferrocyanide, which is white, thrown down. 



170 LESSONS IN CHEMISTRY. 

When potassium ferrocyanide is heated to redness in closed vessels, it is 
converted into potassium cyanide, while iron and carbon separate and nitro- 
gen is disengaged. When it is heated in the air or with certain oxidizing 
agents, it yields potassium cyanate, KOCN. Under the same circumstances 
with sulphur it forms potassium sulphocyanate, KSCN. 

267. Prussian Blue, Ferric Ferrocyanide, (Fe 2 ) 2 (FeC 6 N 6 ) 3 . 
— When ferrous sulphate, FeSO 4 , is added to a solution of potas- 
sium ferrocyanide, the atom of iron changes place with two atoms 
of potassium, and a pale-blue precipitate containing FeK 2 Fe(CN) 6 
is formed. When, however, potassium ferrocyanide is added to a 
ferric salt, such as ferric chloride, FeCP, a dark-blue precipitate 
of Prussian blue is thrown down. As each atom of iron in 
ferric chloride replaces three atoms of hydrogen in as many 
molecules of hydrochloric acid, it will also replace three atoms 
of potassium, and we must write 

4FeCl3 + 3K*Fe(CN)« = 12KC1 + Fe^FeCW) 3 

Ferric chloride. Potassium ferrocyanide. Prussian blue. 

Prussian blue, much used as a pigment, generally comes in 
cubical masses having a coppery reflection. It is insoluble in 
water, and in dilute acids, with the exception of solutions of oxalic 
acid. It is dissolved by alkaline hydroxides, which destroy its color. 

While we are uncertain of the exact relations of the atoms in the mole- 
cules of the ferrocyanides, yet we have learned that they contain a distinct 
radical, ferrocyanogen, Fe(CN) 6 ; and we see that the relations of the atom of 
iron in this radical are quite different from those of the four iron atoms in 
Prussian blue. The latter readily leave and re-enter the molecule by double 
decomposition, but the iron atom in ferrocyanogen always goes with the six 
groups, CN, unless the molecule be decomposed by heat or energetic chemical 
agents. 

268. Potassium Ferricyanide, K 3 Fe(CN) 6 .— This compound 
is formed by passing chlorine gas into a solution of potassium 
ferrocyanide. The chlorine removes one atom of potassium 
from each molecule of the ferrocyanide, thereby converting the 
iron into the ferric state. 

2K*FeC 6 N 6 + CI 2 = 2KC1 + 2K 3 FeC 6 N« 

Potassium ferrocyanide. Potassium ferricyanide. 

The salts are separated by crystallization. 



POTASSIUM ISOCYANATE. 171 

Potassium ferricyanide forms beautiful, large, ruby-red, anhy- 
drous crystals. It dissolves iu about four times its weight of cold 
water, and the solution has a greenish-brown color. It forms no 
precipitate with ferric salts, but with ferrous sulphate gives a dark- 
blue precipitate of ferrous ferricyanide, called Turnbull's blue. 

2K 3 FeC6X6 + 3FeSO* - 3K 2 SO* + Te*(EeC*S*)* 

Potassium ferricyanide. Ferrous sulphate. Turnbull's blue. 



LESSON XXXIII. 



CYANATES AND UREA. 



269. Potassium Cyanate, KO.CN. — When an intimate mix- 
ture of perfectly dry potassium ferrocyanide with half its weight 
of manganese dioxide is heated to dull redness with constant 
stirring, the mixture becomes black and pasty. The potassium 
ferrocyanide has been decomposed, and potassium cyanate ex- 
ists in the product. To extract this substance, the black mass 
is finely powdered and the powder shaken up with boiling eighty 
per cent, alcohol : the liquid is quickly decanted from the sedi- 
ment, and on cooling deposits potassium cyanate in small, color- 
less, anhydrous crystals. It is very soluble in water ; only 
slightly soluble in cold alcohol. When the aqueous solution is 
heated, the cyanate is decomposed into potassium carbonate, 
carbon dioxide, and ammonia. 

2KO.CN + 3H 2 - K2C0 3 + CO 2 + 2XH 3 

Potassium cyanate is decomposed in the same manner by 
acids : hydrochloric acid converts it into potassium chloride and 
ammonium chloride, while carbon dioxide escapes with efferves- 
cence. 

KO.CN + 2HC1 + H 2 = KC1 + NH*C1 + CO 2 

270. The acid corresponding to potassium cyanate is of course 
cyanic acid, HO.CN, but it cannot be made by double decom- 



172 LESSONS IN CHEMISTRY. 

position with potassium cyanate. It has been obtained, how- 
ever, by distilling cyanuric acid* N 3 C 3 3 H 3 . 

N 3 C 3 (OH) 3 = 3H0.CN 

Cyanic acid is very unstable : at ordinary temperatures it rapidly 
changes into an amorphous substance called cyamelide. 
The most interesting salt of cyanic acid is 

271. Ammonium Cyanate, NH 4 O.CN, which is formed when 
vapor of cyanic acid is mixed with ammonia gas. It is a white 
solid, very soluble in water. When its aqueous solution is boiled, 
or even left to itself for a few days, the ammonium cyanate is 
converted into another substance of the same molecular compo- 
sition, CON 2 H*, called urea. 

272. In order to explain this fact, that two substances may 
be represented by the same formula and yet have entirely dif- 
ferent properties, we must believe that the atoms are differ- 
ently arranged within their molecules. Such compounds are 
said to be isomeric, and to possess different molecular structure. 
The structural formulae we employ are deduced from the 
modes of formation, and the decompositions of the compounds 
they represent. Isomeric bodies are very common among the 
carbon compounds. 

273. Urea, CO(NH 2 ) 2 , may be formed by a reaction which 
establishes its molecular structure beyond doubt. When car- 
bonyl chloride, C0C1 2 , is made to react with ammonia, urea and 
hydrochloric acid are formed. 

COCl 2 + 2NH 3 = CO(NH2)2 + 2HC1 
Carbonyl chloride. Urea. 

Here two molecules of ammonia lose each one atom of hydro- 
gen, which combines with the chlorine of the carboDyl chloride, 



* Cyanuric acid is a white crystalline body which is formed by heating urea, 
and by the action of water on the solid chloride of cynogen, C 3 N 3 C1 3 . This 
latter results from the action of chlorine upon hydrocyanic acid in direct sun- 
light. 

3HCN + 3C1 2 = C 3 N 3 C1 3 + 3HC1 



AMMONIUM ISOCYANATE. 1<3 

and the unsatisfied groups CO and 2NH 2 combine, forming a 
molecule of urea. The group NH 2 passes readily from one mole- 
cule to another by double decomposition. It represents a molecule 
of ammonia from which an atom of hydrogen has been removed : 
it is a mouatomic radical. We may then consider that urea is 
formed from two molecules of ammonia by the replacement of 
one atom of hydrogen of each by the diatomic radical carbonyl 
Compounds formed by the replacement of the hydrogen atoms of 
ammonia by other atoms or groups are called amines or amides : 
when the replacement is by the radicals of acids, the name amide 
is used to designate the new compound, while amine is applied to 
such compounds as result from the replacement of these hydrogen 
atoms by radicals which are also capable of replacing the hydrogen 
of acids. Since carbonyl is the radical of carbonic acid, which 
is carbonyl dihydrate, CO(OH) 2 , we call urea carbonyl amide, or 
carbamide. 

274. We have already learned that urea is formed by a curious 
change which takes place in ammonium cyanate. Since we can 
readily prepare potassium cyanate, we have a ready means of 
obtaining urea. For this purpose potassium cyanate is pre- 
pared as has already been described (§ 269) ; but, instead of 
exhausting the mass with alcohol, we exhaust it with cold water, 
which dissolves out the cyanate. The solution is then mixed 
with ammonium sulphate in quantity equal to five-sevenths of 
the potassium ferrocyanide used, and the whole is evaporated to 
dryness on a water-bath. The ammonium sulphate reacts with 
the potassium cyanate, forming potassium sulphate and ammo- 
nium cyanate, and the latter becomes converted into the isomeric 
compound, urea. The mixture of the two bodies is extracted 
with a small quantity of boiling alcohol, which does not dissolve 
the potassium sulphate, but dissolves the urea, and on cooling 
deposits it in crystals. 

The formation of urea from ammonium cyanate was discovered by TVoehler 
in 1828. It was the first instance of the production of an " organic" com- 
pound from a body of mineral origin by chemical means. Before that time 
it was held that compounds found in plants and animals — organic compounds 
— could not be formed except under the influence of a vital force. 



174 LESSONS IN CHEMISTRY. 

275. Urea is the principal solid constituent of the urine : it 
is in this compound that the greater part of the nitrogen of 
burned tissues is removed from the body. It may be extracted 
from urine by evaporating the liquid to a thick syrup, and adding 
nitric acid when it has cooled. The nitric acid combines with 
the urea, forming urea nitrate, CO(NH 2 ) 2 .HN0 3 , which separates 
in a mass of crystals. These are drained, and treated with 
barium carbonate as long as there is effervescence. The mix- 
ture is then evaporated to dryness, and the urea is dissolved 
from the barium nitrate by boiling alcohol. 

276. Urea forms colorless crystals having a cooling taste. It 
dissolves in its own weight of water, and in five times its weight 
of cold alcohol ; it is very soluble in boiling alcohol. An aqueous 
solution of chlorine instantly decomposes it, setting free nitrogen 
and carbon dioxide. 

CO(NH2)2 + H20 + 3C12 = CO 2 + N* + 6HCI 

By the action of heat, its solution in water is converted into 
ammonium carbonate. 

CO(NH2)2 + 2H 2 = (NH^CO 3 

The same reaction takes place slowly in urine, and accounts for 
the ammoniacal odor of stale urine. 

277. Potassium Thiocyanate or Sulphocyanate, KS.CN. — 
A mixture of potassium ferrocyanide with half its weight of 
flowers of sulphur is heated to dull redness in a covered cruci- 
ble. After cooling, the mass is dissolved in water, the liquid is 
filtered, and potassium carbonate is added as Ions: as it causes 
any precipitate. Then the liquid is again filtered, and the solu- 
tion evaporated to dryness. The residue is extracted with hot 
alcohol, and the alcoholic solution allowed to evaporate. Potas- 
sium thiocyanate then separates in colorless, deliquescent crys- 
tals which are very soluble in water and in alcohol. A solution 
of potassium sulphocyanate produces a blood-red color (ferric 
sulphocyanate) with solutions containing ferric salts. 

FeCl 3 + 3KS.CN = Fe(S.CN) 3 + 3KC1 

With silver nitrate it gives a white curdy precipitate which 
is insoluble in nitric acid. 



METHANE. 



175 



Potassium thiocyanate corresponds to the cyanate in which 

the oxygen atom is replaced by an atom of sulphur. 

278. Ammonium Thiocyanate, (NH 4 )S.CX, is found in small quantity in 
the water which has been used to wash coal-gas (§ 225). Representing am- 
monium cyanate in which the oxygen is replaced by sulphur, it undergoes by 
the action of heat a similar curious change into the isomeric compound thio- 
urea, CS(NH 2 ) 2 , whose molecule is exactly like that of urea, excepting that it 
contains sulphur instead of oxygen. 



LESSON XXXIV. 

COMPOUNDS OF CARBON AND HYDROGEN (i). 

279. Methane, CH 4 . — In a glass flask on a sand-bath we heat 
a mixture of equal parts of dried sodium acetate, sodium hydroxide, 
and powdered lime (Fig. 84). The lime does not enter into the 




Fig. 84. 

reaction which takes place, but it prevents the hot sodium hydrox- 
ide from melting through the glass. Since gas will be disengaged, 
we have adapted to our flask a cork and tube, and may collect this 
gas over water, in which it is almost insoluble. The gas is methane : 
it is produced by a reaction between the sodium acetate and sodium 
hydrate, which at the same time yield sodium carbonate. 

XaC2H30 2 + XaOH = Na^CO 3 + CH± 

Sodium acetate. Sodium carbonate. Methane. 



176 LESSONS IN CHEMISTRY. 

280. It is a colorless gas, having no odor. Its density compared 
to air is 0.559, or compared to hydrogen, 8 : this corresponds to a 
molecular weight of 16, as is indicated by the formula, CH 4 . It 
is a combustible gas, and burns with a pale flame. It forms an 
explosive mixture with air or oxygen, and this mixture is often 
unfortunately formed in the galleries of coal-mines, for methane is 
the fire-damp of the miners. It exists under strong pressure in 
the coal-beds, and escapes when these beds are cut into by the 
miners. 

We have already learned that a certain temperature is necessary 
for combustion, as indeed for all chemical action, and a gas cannot 
continue burning when its flame is cooled below the igniting point. 
When a flame is inserted in a tube, not too wide, it is extinguished, 
because the walls of the tube cool it. For this reason the flame 
does not run down the tube of a good Bunsen burner, although 
the combustible gas is mixed with air. A piece of wire gauze 
may be regarded as composed of a large number of fine, short 
tubes, and wire gauze will prevent the passage of flame. The 
fineness of the gauze required will depend on the igniting point 
of the gas or vapor, and, as this temperature is lower, the gauze 
must be finer. We may depress a piece of wire gauze in the flame 
of a Bunsen burner or a lamp, and the flame is kept below the 
gauze until the latter is heated to the temperature required for 

the combustion of the gas. 
Yet the combustible gas 
passes through, and we may 
light it above the gauze : in 
the same manner we may 
hold the gauze a short dis- 
tance above the burner in 
the escaping but unlighted 
gas, and we may ignite the 
gas above the gauze ; the 
flame does not, however, pass below until the gauze becomes 
heated as before (Fig. 85). These principles are applied in the 
miners' safety-lamp, which is practically a lamp so arranged that air 




METHANE. 



177 



can pass to the flame and the burned gases escape only through 
the meshes of fine wire gauze (Fig. 86). For 
better illumination, that part of the gauze im- 
mediately around the flame is usually replaced 
by thick glass. The explosive gases may enter 
this lamp, and may burn inside, but the flame 
cannot pass through unless the gauze become 
highly heated. In most countries it is unlaw- 
ful to continue working galleries containing 
explosive gases until those gases are removed 
by ventilation. The safety-lamp affords a 
means of detecting the presence of very small 
quantities of such gases without danger of ex- 
ploding them. We pass a little illuminating 
gas. of which about 40 per cent, is methane, 
into an inverted jar, and mix it well with the 
air in the jar by moving a roll of paper around 
in it. We now push up into the jar a lighted 
wax taper, the end of which projects just be- 
yond a small glass tube slipped over it, so that 
the flame is quite small. We see that this small 
flame is surmounted by a pale and tremulous 
bluish cap (Fig. 87) : this is owing to the 
combustion of the mixture of gas and air im- 
mediately around the flame, but there is so 
little of the combustible gas present that the 
heat produced by its combustion immediately around the flame is 
not sufficient to carry the combustion throughout the whole mix- 
ture ; otherwise there would be an explosion. By looking at the 
flame in his safety-lamp, the miner can tell by the presence or 
absence of this bluish cap whether any fire-damp be present, and, 
if so, whether there be sufficient to indicate danger of explosion. 
281. Methane is one of the products of the putrefaction of 
vegetable matters in presence of water. It is formed by the de- 
composition of such substances in the muddy bottoms of ponds 
and rivers, and rises in bubbles through the water when this mud 

12 




Fig. 86. 



178 



LESSONS IN CHEMISTRY. 




Fig. 87. 



is stirred : it often collects under the ice in winter, and will escape 
and burn with a pale flame when the ice is pierced and the gas 
lighted. Because of its formation in these 
localities, methane is often called marsh gas. 

282. The composition of methane shows us that the 
carbon atom is tetratomic ; it has the combining power 
of four atoms of hydrogen. We have already learned 
that chlorine has an energetic affinity for hydrogen, 
and that it will remove this element from many hy- 
drogen compounds. When chlorine is mixed with 
methane, and the mixture is exposed to light, the 
chlorine removes the hydrogen from the methane, and 
hydrochloric acid is formed, but an atom of chlorine 
takes the place of every atom of hydrogen so removed. 
We may consider that there is a double decomposition 
between the chlorine molecules and the methane mole- 
cules, and this decomposition may continue until all 
the hydrogen atoms of the methane are replaced by chlorine. 

CH± + CI 2 = CH 3 C1 + HC1 

CH* + 2C1 2 = CH 2 C1 2 + 2HC1 

CH 4 + 3C1 2 = CHC1 3 + 3HC1 

CH* + 4C1 2 = CC1 4 + 4HC1 

All these compounds of carbon with chlorine and hydrogen may thus be ob- 
tained by substitution. Their compositions are a still further evidence that 
the carbon atom is tetratomic. The hydrogen atoms of methane may also be 
replaced by the monatomic atoms of bromine and iodine, producing compounds 
precisely similar to those formed by chlorine. 

One of the substances so formed has the composition CH 3 I; it is called 
methyl iodide, and the compound CH 3 C1 is called methyl chloride. We may 
consider that the group of atoms CH 3 acts like a single atom as potassium in 
potassium chloride ; and when we have learned that it may take part in double 
decompositions, leaving one molecule and entering another without change, 
we shall see that it is a radical; it is called methyl. 

283. Methyl iodide, CH 3 I, is a colorless liquid. When it is 
sealed up in strong glass tubes containing some zinc, and the 
tubes are heated for a time to about 150°, the zinc takes away 
the iodine from the methyl iodide, and zinc iodide, Znl 2 , is 
formed. When the tubes are carefully opened, they are found to 
contain a gas to which both analysis and density assign the com- 
position C 2 H 6 . How must the atoms be related in a molecule of 
this gas ? Are the carbon atoms still tetratomic ? How has the 



COMPOSITION OF HYDROCARBONS. 179 

gas been formed ? We must believe that when two atoms of 
iodine are removed from two molecules of methyl iodide the two 
monatomic methyl groups, CH 3 , combine together ; that in a mol- 
ecule of the gas, C 2 H 6 , the two carbon atoms, each with its three 
hydrogen atoms, like three satellites, form a perfect system. We 
can represent this relation by our formulae. 

CH 3 I + Zn + ICH 3 - Znl 2 + H 3 C-CH 3 

Methyl iodide. Zinc. Methyl iodide. Zinc iodide. Ethane. 

Then in this gas, C 2 H 6 , which is called ethane, the affinity of 
the carbon atoms must be satisfied partly by their combination 
together, and partly by their combination with hydrogen. 

284. By the action of chlorine on ethane, the hydrogen of that 
gas may be replaced by chlorine atoms, and compounds may also 
be obtained in which the replacement is by iodine atoms. When 
only one of the hydrogen atoms is so replaced, the compound 
C 2 H 5 I is formed. We consider that this contains the radical 
C 2 H 5 , which is called ethyl, and the molecule C 2 H 5 I is called 
ethyl iodide. Since all the atoms of hydrogen in a molecule of 
ethane must have the same relations to the carbon atom around 
which they move, and also to the other carbon atom, it is a matter 
of indifference which one we suppose to be replaced by the iodine 
atom. When ethyl iodide and methyl iodide, in the proportions 
required for the same number of molecules of each, are heated 
with zinc in sealed tubes, a reaction takes place just as in the case 
of zinc and methyl iodide alone. That is, both iodine atoms are 
removed, and we may say either that the iodine of ethyl iodide 
is replaced by the group methyl, CH 3 , or that the iodine of methyl 
iodide is replaced by the radical ethyl, C 2 H 5 . A gas called pro- 
pane, C 3 H 8 , is then formed. 

C 2 H5I + Zn + ICH 3 = Znl* ' + C 2 H5-CH 3 
Ethyl iodide. Methyl iodide. Zinc iodide. Propane. 

We find, then, that the atoms of carbon are able to combine to- 
gether; that they form complex systems in which each carbon atom 
is accompanied by atoms of hydrogen or some other element. 
As we have done before, we may compare the carbon atoms to 
stars or suns which revolve around each other ; each sun is ac- 



180 LESSONS IN CHEMISTRY. 

companied by its own planets, and we shall presently see that 
each of the planets may have its satellites. 

By reason of the property of combination between its own 
atoms, a property which is not possessed in the same degree by 
the atoms of any other element, carbon forms an almost infinite 
number of compounds. These compounds differ from those of 
the other elements in this respect : — while any other element forms 
a few compounds with nearly all other elements, carbon forms in- 
numerable compounds containing very few of the other elements. 
The more numerous of the carbon compounds contain only carbon, 
hydrogen, oxygen, and nitrogen, but all of the other elements 
may, under proper conditions, be made to form part of these com- 
pounds. The carbon compounds are generally called organic 
compounds. 

285. The compounds containing carbon and hydrogen only, are 
Called hydrocarbons ; we have just studied three of them, and in 
the molecules of each of these the combining power of the car- 
bon atoms is completely exhausted. We may express in detail 
the atomic relations of the three. 



H 


HH 


HHH 


H-C-H 


H-C-C-H 


H-C-C-C-H 


H 


H H 


HHH 


Methane. 


Ethane. 


Propane. 



The union of the carbon atoms together does not stop with propane, for in 
turn one of its hydrogen atoms may be replaced by a methyl group, and the 
hydrocarbon, C 4 H 10 , is the result. In the same manner this may be converted 
into C 5 H 12 , and a whole series of saturated hydrocarbons has been obtained. 
When we examine the composition of the members of this series, we see that 
each contains two more than twice as many atoms of hydrogen as it does of 
carbon. We may express the composition of any member of the series by the 
general formula C n H 2n + 2 , n representing the number of carbon atoms in the 
molecule. The names of these compounds end in ane, and after the fourth 
member, the prefix indicates the number of carbon atoms in a molecule. 

CH*, Methane. C^H 12 , Pentane. 

C 2 H6, Ethane. C 6 H 14 , Hexane. 

C 3 H8, Propane. C 7 H 16 , Heptane. 

C*Hi°, Butane. C 8 H!8, Octane. 

The first five are gases at ordinary temperatures ; the others are liquids of 
which the boiling points are higher as the number of carbon atom3 in the 



PETROLEUM. 181 

molecule increases, until, when this number reaches sixteen, the compounds 
are solid at ordinary temperatures. Ordinary paraffin is a mixture of the solid 
members of the series; its name, meaning poor affinity, indicates that it does 
not readily enter into chemical reactions, and, since this property is common 
to all of the saturated hydrocarbons, the series C n H 2n + 2 is often called the 
paraffin series. 

We see that each member of this series contains one atom of carbon and two 
atoms of hydrogen more than the preceding. Compounds which thus differ 
from each other by CH 2 , or a multiple of that symbol, and which have the 
same general chemical properties, are said to be homologous, and to form a 
homologous series. 



LESSON XXXV. 
COMPOUNDS OF CARBON AND HYDROGEN (2). 

286. Petroleum. — Petroleum, or rock-oil, as the name signi- 
fies, has been known from very early history, but it has been 
marvellously abundant in commerce only since 1859, when it was 
found that the oil would flow from wells bored into the rock in 
Northwestern Pennsylvania. The oil usually occurs in a loose, 
coarse sandstone into which it has drained from its source in other 
rocks. That source is still a matter of uncertainty, but the oil 
has doubtless been formed by the decomposition of vegetable and 
perhaps animal matters, long buried in the earth. The depth to 
which the wells must be sunk varies with each locality ; some- 
times it is only a few feet ; sometimes it may be two or three 
thousand feet. Sometimes the oil begins to flow as soon as the 
oil-bearing rock is penetrated, but more usually the interior press- 
ure is not strong enough to raise the oil, and a pump must then be 
employed. Petroleum is widely distributed, being met with in 
nearly all parts of the globe. The principal localities are in the 
Eastern United States and Canada, Russia, Austria, and Eastern 
Asia. 

Crude petroleum varies in color from pale yellow to almost 
black ; it usually has a greenish tint. It is sometimes quite fluid, 
sometimes thick like molasses. Its density is comprised between 
0.75 and 0.92. It is a mixture of a large number of hydrocarbons, 



182 LESSONS IN CHEMISTRY. 

of which those found in American oil chiefly belong to the par- 
affins which we have just studied. Indeed, all the saturated 
hydrocarbons, from CH 4 up to C 16 H 34 , have been separated from 
it ; the hydrocarbons of the Russian oil belong to a different 
series. Crude petroleum is not used for illuminating purposes, 
but is refined by a process called fractional distillation. This 
consists in slowly heating the oil and collecting separately the 
portions that pass off at different temperatures. That which 
distils over below 70° is called naphtha ; the temperature is then 
raised to about 150°, and the liquid condensed up to that point 
is benzine : between 150° and 280°, kerosene, or illuminating oil, 
distils over, and that portion which passes between 280° and 400° 
is paraffin oil or lubricating oil. Much paraffin distils towards 
the close of the operation, and a residue of coke remains in the 
retort. 

Naphtha has a density of about 0.65, and, when purified from 
its most volatile constituents, forms gasoline, used in some gas- 
machines. Air is blown through the gasoline, and becomes charged 
with sufficient of the vapor of the volatile hydrocarbons to burn 
with an illuminating flame. Benzine has a density of about 0.702, 
and boils at about 148°. It is used for dissolving oils and fats, and 
instead of turpentine for mixing with paints. 

Kerosene should contain no product whose boiling point is 
below 150°, for the vapors of the more volatile hydrocarbons form 
dangerously explosive mixtures with air. The fire-test by which 
the safety of the oil is determined, is made by slowly heating the 
oil in a little dish on a water-bath, carefully observing by means 
of a thermometer the temperature at which inflammable vapors 
are given off and the temperature at which the oil takes fire. 
A lighted match is passed rapidly over the oil, about half a cen- 
timetre from its surface ; when the vapor burns with a little 
flash, the thermometer marks the flashing point. A few degrees 
above this, the oil itself takes fire. The flashing-point should not 
be below 60°, and the burning-point not below 65°. 

287. Paraffin. — The name paraffin is commonly applied to 
that product of the distillation of petroleum which solidifies on 
cooling : it is also a product of the destructive distillation of peat 



UNSATURATED HYDROCARBONS. 183 

and some kinds of coal. When the last liquid portions of the dis- 
tillate of petroleum are cooled by ice, a considerable quantity of 
paraffin separates. When purified, paraffin is a colorless, trans- 
parent, or translucent mass. It is a mixture of several members 
of the series of saturated hydrocarbons. Accordingly as it has 
been prepared and purified, its melting-point varies from 45° to 
65°. It makes excellent candles. 

288. We have now learned something about one class of hydrocarbons, a 
class in which the carbon atoms cannot combine with any other atoms unless 
they separate from each other. It is worthy of notice that they do not sepa- 
rate from each other except by the action of the most energetic agents : on the 
contrary, these carbon atoms remain combined, and, with as many of their 
accompanying hydrogen atoms as we permit to remain with them, constitute 
fixed and definite radicals, which act exactly like the atoms of elements having 
the same combining powers. We have noticed two of these radicals, methyl 
and ethyl. In order that there may be a uniformity of names for these com- 
plex groups, chemists have agreed to retain the first syllable of the name of 
the saturated hydrocarbon, in the names of all compounds derived from that 
hydrocarbon. The termination in yl has been selected for the radicals which 
we consider are formed by the removal of one atom of hydrogen from a satu- 
rated hydrocarbon, and then the first word of the name of a compound will 
show us the hydrocarbon radical in the molecule, and the last word must indi- 
cate the atom or group of atoms combined with that radical. We may then 
understand the composition of the following bodies : 

CH±, Methane. CH 3 Br, Methyl bromide. CH 3 .OH, Methyl hydroxide. 
C 2 H6, Ethane. C 2 H 5 C1, Ethyl chloride. C 2 H5.0H, Ethyl hydroxide. 

C 3 H 8 , Propane. C 3 H 7 NH 2 , Propyl amine. (C 3 IF) 2 0, Propyl oxide. 

When we have once acquired definite ideas of what is meant by a radical, 
that it is a group which acts precisely as an atom, having continually the same 
combining power or atomicity, leaving one molecule and entering another as a 
distinct existence; then the structure of these complex molecules becomes per- 
fectly intelligible, and we need only be acquainted with the radicals concerned 
in order to be able at once to interpret, by our system of atomic groupings, the 
relations of the atoms in the molecule of any compound. 

289. Unsaturated Hydrocarbons. — We have mixed in a glass 
flask some alcohol with four times its weight of strong sulphuric 
acid, and, as this mixture sometimes froths very much when we 
heat it, we have put in enough sand to absorb the liquid almost 
entirely. After fitting to our flask a cork through which pass a 
delivery-tube, and a safety-tube in which we put a little mercury 



184 LESSONS IN CHEMISTRY. 

or some sulphuric acid, we heat it on a sand-bath. A gas is dis- 
engaged, and we may collect it in jars over the pneumatic trough. 

290. Ethylene, C 2 H 4 . — The gas which we have prepared is a 
hydrocarbon. It is colorless and almost odorless : its density com- 
pared to air is 0.9784, or compared to hydrogen, 14. Analysis 
shows that it contains carbon and hydrogen in the proportion of 
one atom of the first to two atoms of the second, and its density 
shows that its molecule must contain two atoms of carbon and four 
of hydrogen. Its composition is, then, C 2 H 4 : it is called ethylene. 
It burns with a brilliant flame. 

Into a jar of this gas we pour a little bromine, and cause it to 
flow over the sides of the jar : the color of the bromine disappears, 
and drops of an oily liquid are formed. This liquid has a pleasant 
odor, very different from the suffocating vapor of the bromine. 
The ethylene has combined with the bromine and formed this 
liquid, which is called ethylene bromide. The vapor-density and 
analysis of the compound assign to its molecule the composition 
C 2 H 4 Br 2 . Evidently if the molecule C 2 H 4 can combine directly 
with two atoms of bromine, it must be a diatomic molecule, capable 
of manifesting the combining power of two atoms of hydrogen. 
Let us study the reaction by which ethylene is formed : alcohol is 
ethyl hydrate, C 2 H 5 .OH : sulphuric acid, by its strong affinity for 
water, converts it into H 2 -|- C 2 H 4 . Then, in losing the mon- 
atomic hydroxyl group and an atom of hydrogen, the carbon atoms 
of alcohol must recover the combining powers of two atoms of 
hydrogen : this combining power is manifested in the combination 
of ethylene with bromine, chlorine, etc. If two atoms of hydrogen 
were removed from a molecule of methane, CH*, the remaining 
group, CH 2 , would be diatomic, and we believe that a molecule of 
ethylene gas is formed by the union of two such diatomic groups, 
and is expressed by the formula CH 2 -CH 2 ; but these atoms 
then possess more energy than when combined in the gas ethane. 
CH 3 -CH 3 , and may develop that energy and enter into direct com- 
bination with bromine, forming ethylene bromide, CH 2 Br-CH 2 Br. 

Ethylene Chloride, CH 2 C1-CH 2 C1, is formed when equal 
volumes of chlorine and ethylene are mixed in diffuse daylight. 



DIATOMIC HYDROCARBONS. 185 

It is a somewhat oily liquid, and from this character ethylene was 
first called olefiant gas. It boils at 82°. 

Ethlyene Bromide, CH 2 Br-CH 2 Br, is made by passing ethy- 
lene gas into cooled bromine. It boils at 131°. 

291. We have seen that chlorine is capable of replacing the 
hydrogen of ethane, C 2 H 6 , atom for atom. From the products of 
this reaction we can by careful operations separate two liquids 
having the composition C 2 H 4 C1 2 , but having entirely different 
properties. These compounds are isomeric, and we may understand 
their isomerism when we see that both atoms of chlorine may 
replace hydrogen atoms which are in relation to the same atom 
of carbon, forming the molecule CH 3 -CHC1 2 ; or each may replace 
an atom of hydrogen from a group CH 3 ; the compound formed 
in the latter case would of course be ethylene chloride. 

292. Ethylene is only the first member of a long series of hydrocarbons 
which we may consider are derived from it by the replacement of one or more 
of its hydrogen atoms by the monatomic hydrocarbon radicals which we have 
already studied. Each of the compounds so formed is diatomic : it will com- 
bine directly with two atoms of chlorine or bromine, and may be made to 
combine with two monatomic radicals or with one diatomic radical. The names 
of these diatomic hydrocarbons are made to correspond with the saturated 
hydrocarbons, from which we may consider they are derived by the removal 
of two atoms of hydrogen, but the ane of the name is changed to ylene. Ethy- 
lene corresponds to ethane, butylene corresponds to butane. We have here 
our second series of homologous compounds, each differing from the next by 
CH2. 

C 2 H±, Ethylene. C 5 H 10 , Amylene or pentylene. 

C 3 H 6 , Propylene. C 6 H 12 , Hexylene. 

C 4 H 8 , Butylene. C 7 H U , Heptylene, etc. 

On examination, we notice that each molecule contains twice as many 
atoms of hydrogen as of carbon ; the general formula for the series is C n H 2n . 
The proportion of hydrogen and carbon is the same in each member of the 
series, but the molecular weights, and consequently the number of atoms in 
the molecules, are not the same. Bodies of which the molecules contain the 
same atoms in the same proportion but in different numbers are said to be 
polymeric. All of these diatomic hydrocarbons are polymeric ; the number 
of carbon atoms and hydrogen atoms in each is an exact multiple of CH 2 . 
Because these hydrocarbons combine directly with chlorine and bromine, 
forming oily liquids, the series is often called the olejine series. 

293. It has been said that we may consider these bodies as formed from 
ethylene by the replacement of hydrogen atoms by the monatomic radicals, 



186 LESSONS IN CHEMISTRY. 

methyl, ethyl, etc. We must see that this replacement may yield many in- 
stances of isomerism. If one of the hydrogen atoms of ethylene be replaced 
by methyl, we obtain propylene. 

CH 2 =CH2 CH 2 --CH-CH 3 

Ethylene. Propylene. 

By the replacement of two of the hydrogen atoms by methyl, we may 
obtain two different butylenes, according to the positions of the replaced hydro- 
gen atoms, and there is still a third butylene, formed by the replacement of 
one hydrogen atom by an ethyl group. 

CH2 CH.CH 3 C(CH 3 ) 2 CH(C 2 H5) 

CH2 CH.CH 3 CH2 CH2 

Ethylene, (a) Dimethylethylene. (|8) Dimethylethylene. Ethylethylene. 
All these hydrocarbons have been obtained and studied, and their names 
indicate the molecules from which they are derived and the radicals which 
are substituted for the hydrogen atoms in those molecules. 

Acetylene, C 2 H 2 , is the first member of another series of 
unsaturated hydrocarbons. We may prepare it by pouring 
water over calcium carbide, a crystalline body which is made 
from lime and coke in the electrical furnace. 

CaC 2 + 2H 2 = C 2 H 2 + Ca(OH) 2 

Acetylene is a colorless gas having a characteristic odor. It 
is slightly soluble in water, and liquefies under great pressure. 
It burns with a very smoky, luminous flame. With air or oxy- 
gen it forms a highly explosive mixture. Acetylene has lately 
been introduced as an illuminant : it can be made to yield a 
light of dazzling brilliancy. 



LESSON XXXVI. 

COMPOUNDS OF CARBON AND HYDROGEN (3).— 
ANALYSIS OF CARBON COMPOUNDS. 

294. The tar which condenses during the distillation of bitu- 
minous coal for the manufacture of gas, is an exceedingly com- 
plex liquid, consisting principally of compounds of carbon and 
hydrogen. Some of these compounds are solid, some of them are 
volatile liquids. Since they boil at different temperatures, they 



BENZENE. 



187 




can be separated from each other by fractional distillation. The 
vapors of the substances are passed through a tube which is 
maintained at the boiling point of the most volatile constituent 
of the mixture : in this tube the liquids having higher boiling 
points are condensed, and flow back into the still, while the 
vapor of the more volatile liquid passes on and is condensed 
separately. A simple laboratory contrivance 
is a rather wide tube on which a couple of 
bulbs are blown ; this is placed vertically 
in the flask in which we boil the mixed 
liquid (Fig. 88). The lower part of the 
tube becomes heated to the temperature of 
the mixed vapor, which is between the boil- 
ing points of the liquids : as some of the 
most easily condensed vapor is cooled and 
converted into a liquid, the temperature of 
the tube gradually falls towards the upper 
portion, and by carefully regulating the boil- 
ing, only the most volatile liquid passes from 
the apparatus. This is indicated by a ther- 
mometer of which the bulb is opposite the 
side-tube. 

295. Benzene, C 6 H 6 .— The most volatile 
constituent of coal-tar is a liquid called ben- 
zene. It freezes at 5.5°, and boils at 80.5°. 
It does not dissolve in water, but is soluble 
in alcohol and ether. It is very inflammable, and burns with 
a bright but smoky flame. The composition of its molecule 
is C 6 H 6 , and yet in most of its reactions it acts like a satu- 
rated hydrocarbon. We put a few crystals of iodine into some 
benzene in a glass flask, and pass chlorine through the liquid ; 
hydrochloric acid gas is given off, and the hydrogen atoms of 
the benzene are replaced by chlorine. 

C 6 H6 + CP = C 6 H5C1 + HCl 

The iodine only helps to break up the molecules of chlorine. 



Fig. 



188 



LESSONS IN CHEMISTRY. 



HC 
HC,<^CH 



HC 



CH 



CH 



Evidently the molecular structure of benzene must be different from that of 
the other hydrocarbons which we have studied, and we can only account for its 
resemblance to the saturated hydrocarbons by supposing that its carbon atoms 
are differently related. Of several theories which have been proposed in order 
to explain the chemical behavior of benzene, we need only con- 
sider one, which supposes that the atoms form a complex sys- 
tem, in which each carbon atom is combined with two other 
carbon atoms, and with an atom of hydrogen. The six car- 
bon atoms thus form a closed chain, and we may represent this 
by a hexagon, a carbon atom being placed at each angle. The 
fourth atomicity of each carbon atom may be supposed to form 
a double bond with one of the adjacent carbon atoms. 

296. All the hydrogen atoms of benzene may be replaced by 
other atoms or radicals, and, when more than one is so replaced, 
we have interesting isomeric compounds, the isomerism depending 
on the relations of the carbon atoms whose hydrogen atoms are 
aifected. Let us suppose, for example, that two hydrogen atoms 
are replaced by two chlorine atoms : experiment has shown that 
three compounds may then be formed, having precisely the same 
composition, but different properties. 

We can interpret this by our theory and our representation of the molecule, 
by considering that while one atom of chlorine always occupies the same place, 
the position of the other varies. 

CC1 CC1 CC1 



HC 



HC 



CC1 



CH 



HC 



HC 



CH 



CC1 



HC 



HC 



CH 



CH 



CH 



CH 



CC1 



These formulae represent the three cases in which two atoms of hydrogen in 
benzene are substituted by two atoms of chlorine. Theory does not indicate 
the existence of other isomers. By different methods chemists have always 
succeeded in producing three isomeric compounds in which two atoms of hy- 
drogen of benzene are replaced by other atoms or radicals, but they have 
never been able to obtain more than three such compounds. 

297. There are many hydrocarbons which we believe to be 
derived from benzene by the replacement of its hydrogen atoms 
by radicals, such as methyl, ethyl, etc. Some of these have been 
obtained by methods which allow no doubt as to their constitu- 
tion ; others have not yet been so formed, but certain of their 
chemical reactions seem to show that they also are derived from 
benzene. These hydrocarbons and many of the bodies derived 



TURPENTINE. — NAPHTHALENE. 189 

from them have peculiar aromatic odors, and for this reason the 
whole series of compounds which are considered as benzene de- 
rivatives is commonly called the aromatic series. Of these com- 
pounds we can consider only a few. 

298. Methyl-benzene, or Toluene, C 6 H 5 .CH 3 , was first derived 
from tolu balsam. It is now obtained from coal-tar, and consti- 
tutes a considerable proportion of the benzene of commerce. It 
resembles benzene, but boils at 110°, and does not solidify even 
at —30°. 

We can understand that there may be four isomeric compounds formed by 
the replacement of a single hydrogen atom of methyl-benzene, for that replace- 
ment may affect either a hydrogen atom in one of three places in the benzene 
group, C 6 H 5 , or a hydrogen atom of the radical methyl, CH 3 . 

The three isomeric dimethyl benzenes, C 6 H 4 (CH 3 ) 2 , are called xylenes : two are 
liquids, and one is a solid. Isomeric with them is also ethyl-benzene, C 6 H 5 .C 2 H 5 . 

299. Oil of Turpentine, C 10 H 16 , is a derivative of benzene, and 
it is isomeric with a large number of essential oils. The oils of 
lemon, orange, bergamot, juniper, lavender, and many others, all 
appear to have the same molecular composition, and we must 
believe that their differences are due to a different arrangement 
of the atoms constituting their molecules. These oils are obtained 
by distilling with water the leaves or other parts of the plant con- 
taining them. It is true that the boiling point of each of these oils 
is much higher than that of water, but the steam of the water 
readily carries over the oil. The condensed liquid then sepa- 
rates into two layers, the lower being water, and the upper and 
lighter being the essential oil. Oil of turpentine is so made by 
distilling with water the crude turpentine which flows from in- 
cisions made in certain species of pine-trees. There are several 
varieties of this oil, which differ according to the species of pine- 
tree which furnishes them. The density is about 0.87, and they 
boil at about 156°. Oil of turpentine and most of the essential 
oils slowly absorb oxygen from the air, and are converted into 
various resins. 

300. Naphthalene, C 10 H 8 , is a solid hydrocarbon derived from 
coal-tar. It usually occurs as pearly scales, melting at 79°, and 
boiling at 218°. It does not dissolve in water, and but slightly 
in cold alcohol. It is soluble in boiling alcohol, and crystallizes 



190 LESSONS IN CHEMISTRY. 

when the solution cools. It is employed for the manufacture of 
numerous beautiful dye-stuffs, analogous to the aniline, dyes, which 
we will presently study. 

301. Anthracene, C U H 10 , is one of the least volatile hydro- 
carbons obtained from coal-tar. When pure, it forms beautiful 
transparent prisms, which melt at 213° ; its boiling point is 360°. 
It is employed for the manufacture of alizarin, a red coloring 
matter which was until within a few years obtained only from 
madder. The ability to produce this dye-stuff by purely chemical 
processes has permitted large areas of land which were formerly 
devoted to the cultivation of the madder-plant to be used for 
raising grain. Besides this, chemists have been able to prepare 
from this same alizarin valuable dye-stuffs of other colors, and 
there is now a whole series of anthracene coloring matters. 

ANALYSIS OF CAKBON COMPOUNDS. 

302. The proportions in which the elements exist in any carbon compound 
are determined by elementary analysis. If the compound contains other ele- 
ments than carbon, hydrogen, and oxygen, its analysis requires several opera- 
tions : if only these three elements be present, the carbon and hydrogen are 
determined by one operation, and the quantity of oxygen is the difference be- 
tween the sum of the carbon and hydrogen and the total weight of the substance 
analyzed. The analysis is conducted by mixing a weighed quantity of the 
substance with pure and dry cupric oxide in a long glass tube, one end of 
which is drawn out to a fine point and sealed. The other end is connected 
with a small U-tube containing pumice-stone wet with sulphuric acid, and the 
U-tube is connected with a bulbed tube containing a solution of potassium 
hydroxide (Fig. 89). The tube is heated to redness in a long tube-furnace, and 
the oxygen of the cupric oxide converts the hydrogen of the carbon com- 
pound into water, while the carbon is burned into carbon dioxide. The water 
is absorbed in the tube containing the pumice and sulphuric acid, while the 
carbon dioxide is absorbed by the potassium hydroxide. Towards the close of 
the operation, a caoutchouc tube connected with an oxygen gas-holder is 
slipped over the drawn-out point of the combustion-tube ; the oxygen is turned 
on, and the point is broken off by pinching the end of the tube. A current of 
pure dry oxygen is then passed through the red-hot tube, and all traces of 
carbon dioxide and watery vapor are forced through the absorption-tubes ; at 
the same time any unburned carbon is completely consumed, and the copper 
from which oxygen has been removed is again converted into cupric oxide 
for another operation. The increased weight of the U-tube (j and g) f in which 
water has been absorbed, is due to the water, and one-ninth of the increase 
will represent the quantity of hydrogen in the amount of substance analyzed. 



ANALYSIS OF CARBON COMPOUNDS. 



191 




192 LESSONS IN CHEMISTRY. 

The increased weight of the bulbed tube (h) is due to carbon dioxide, and 
if > or T 3 r , of this increase will give us the quantity of carbon which we wish to 
determine. Since the current of oxygen would carry a little vapor of water 
out of the bulbed tube, and so diminish its weight, a small tube (i) containing 
pumice and sulphuric acid is attached, and in this the vapor is retained; this 
tube is always weighed with the potash bulbs. 

Having determined the proportions of all the elements in a compound, its 
molecular weight is calculated from its vapor-density, or, if this be not pos- 
sible, by other methods. Knowing the molecular weight and the propor- 
tion of each element present, it is very easy to fix the chemical formula 
expressing the composition of the molecule. 



LESSON XXXVII. 

ALCOHOLS (i). 

303. When the iodide of a radical like methyl or ethyl is 
heated with silver oxide and water, silver iodide is formed, and 
the iodine of the carbon compound is replaced by a hydroxyl 
group. 

2CH3I + Ag20 + IPO = 2AgI + 2CH3.0H 
Methyl iodide. Silver oxide. Silver iodide. Methyl hydroxide. 

A hydroxide of the hydrocarbon radical is so formed, and these 
hydroxides constitute what are called the alcohols. We must study 
some of the more important of these compounds. 

304. Methyl Alcohol, CH 3 .OH.— The liquid which condenses 
during the manufacture of charcoal (§ 226) contains small quan- 
tities of a volatile liquid which can be separated by careful frac- 
tional distillation. This liquid is usually sold under the name 
methylene or wood-spirit. It is impure methyl alcohol, and is 
used for the manufacture of varnishes, and for dissolving fats and 
oils. Methyl alcohol has the property of forming with calcium 
chloride a crystalline compound, and is usually purified by satu- 
rating the wood-spirit with calcium chloride, and evaporating the 



ALCOHOLS. 193 

solution by a gentle heat until it crystallizes. The crystals are 
dissolved in water ; when their solution is boiled, the compound 
of methyl hydrate and calcium chloride is decomposed, and the 
methyl alcohol can be separated by fractional distillation. 

Pure methyl alcohol is a colorless liquid, of an odor resembling 
that of common alcohol. Its density at 0° is 0.814, and it boils 
at 66.5°. It mixes in all proportions with water and with ordinary 
alcohol. It is inflammable, and burns with an almost colorless 
flame. We throw a piece of sodium into methyl alcohol : hydrogen 
is given off, and the metal dissolves : an atom of sodium has re- 
placed the hydrogen atom of the group hydroxyl, and sodium 
methylate is formed. The reaction is precisely like that which 
yields sodium hydroxide in the reaction of sodium with water. 
2H-0-H + 2Na = 2NaOH + H2 

2CH 3 -0-H + 2Na = 2CH 3 -ONa + H 2 
Methyl alcohol. Sodium methylate. 

When methyl alcohol is oxidized slowly, an atom of oxygen 

replaces two hydrogen atoms of the methyl group, CH 3 , and formic 

acid results. 

CH3.0H + O 2 = H 2 + CHO.OH 
Methyl hydroxide. Formic acid. 

305. Ethyl Alcohol, C 2 H 5 .OH.— When ethylene gas is passed 
into strong hydriodic acid, direct combination takes place, and 
ethyl iodide is formed. 

C 2 H* + HI = c 2 H5i 
When this ethyl iodide is heated with potassium hydroxide solu- 
tion, a double decomposition takes place, yielding potassium iodide 
and ethyl alcohol. 

C*IM + KOH = KI + C 2 H5.0H 
However, ethyl hydroxide, which is ordinary alcohol, is manufactured 
by a peculiar decomposition of glucose, or some substance having 
the same composition as glucose. This decomposition is brought 
about by a minute organism which lives and multiplies by con- 
verting the glucose into carbon dioxide and alcohol. A decompo- 
sition due to such an organized being is called a fermentation, and 
the organism is called a ferment. The molecule of glucose is 

13 



194 LESSONS IN CHEMISTRY. 

expressed by the formula C 6 H 12 6 , and, although small quantities 
of other substances are produced during the fermentation, which 
is caused by the yeast-plant and is called the alcoholic fermenta- 
tion, the general change may be represented by the equation 
C 6 H 12 06 = 2C 2 H5.0H + 2C0 2 

For the manufacture of alcohol, the product of the fermentation 
is distilled, and the alcohol so separated from the water. However, 
the best apparatus does not give alcohol stronger than about ninety- 
five per cent. Pure or, as it is commonly called, absolute alcohol 
is made by putting quick-lime into the strongest alcohol of com- 
merce, and distilling the mixture after it has stood several days. 
By reason of its strong affinity for water, the lime then retains all 
of that liquid. 

Pure alcohol is a colorless liquid, having a faint but pleasant 
odor. Its density at 0° is 0.8095, and it boils at 78.4°. It mixes 
with water in all proportions, and the mixture becomes slightly 
warm and contracts in volume. Alcohol dissolves many substances 
which are insoluble in water ; among these are iodine, the essential 
oils, fats, and resins. The spirits of the pharmacies, such as spirits 
of ammonia, are solutions of volatile substances in alcohol ; tinc- 
tures are similar solutions of non-volatile substances. 

Alcohol is combustible, and burns with a pale flame, the prod- 
ucts of the combustion being carbon dioxide and water. 

By the slow oxidation of alcohol, acetic acid and a volatile 
liquid called aldehyde are formed ; acetic acid by the replacement 
of two atoms of hydrogen of the ethyl group by an atom of oxy- 
gen, and aldehyde by the replacement of the hydroxy 1 group and 
one atom of hydrogen by an atom of oxygen. Water is of course 
formed in both cases. 

CH 3 -CH 2 .OH f O = CH3-CHO + H 2 

Alcohol. Aldehyde. 

CH 3 -CH 2 .OH + O 2 = CH3-CO.OH + H 2 
Alcohol. Acetic acid. 

The slow oxidation of alcohol may be made to develop consid- 
erable heat. Over a little alcohol in a beaker we suspend a coil 
of platinum wire which we have previously heated to redness 




ALCOHOL. 195 

(Fig. 90). The wire becomes bright red, and will continue to 

glow as long as sufficient air and alcohol vapor come in contact 

with it. At the temperature of the red-hot wire 

the alcohol vapor is fully oxidized, but if we 

remove it, and allow it to cool slightly, and then 

withdraw it before it becomes bright, we may 

notice the peculiar odor developed in the beaker. 

This is due to the formation of aldehyde. 

Fig 90. 
To test for alcohol in solution we add a flake 

or two of iodine, and then potassium hydroxide until the brown 
color disappears. A yellow precipitate of iodoform results if alco- 
hol be present. The smell of this iodoform is very characteristic. 
306. The reaction of alcohol with solutions of certain metals in nitric acid 
yields a class of bodies called the fulminates. Fulminating mercury, which is 
used for charging percussion-caps, may be prepared by dissolving about two 
grammes of mercury in fifteen cubic centimetres of strong nitric acid contained 
in a rather large flask or beaker. The reaction is aided by a gentle heat, and 
as soon as all the mercury has disappeared, the vessel is removed from the 
proximity of flame, and twenty cubic centimetres of alcohol are added. A 
violent reaction takes place, dense, white, poisonous vapors are disengaged, 
and fulminate of mercury is deposited as a light-gray powder. "When the 
effervescence has ceased, the vessel is filled with water, and the acid liquid is 
poured off: the mercuric fulminate is washed by decantation, until the water 
no longer becomes acid. It is then collected on a small filter, and dried by 
exposure to the air. The reaction b} T which this compound is formed is very 
complicated, but the composition of mercuric fulminate is expressed by the 
formula HgC 2 X 2 2 , and its molecule is believed to represent methane, CH 4 , in 
which the hydrogen is replaced by NO 2 , a cyanogen group, and mercury. 
CH 4 , Methane. C(N0 2 )(CX)Hg, Fulminate of mercury. 

Fulminate of mercury explodes violently by friction or percussion, and should 
be kept in loosely- corked bottles, lest it be exploded by the friction of a glass 
stopper. It explodes also at a temperature of about 1S0°. Although this body 
is so exceedingly explosive that it would burst a gun-barrel in which it was 
detonated, the expansive force of the gases produced is much inferior to that of 
those disengaged by gunpowder, and it could not be used for projectile effects. 

307. Alcoholic beverages are products of the fermentation 
of substances containing glucose or some body capable of being 
converted into glucose. In the manufacture of wine, the glu- 
cose is derived from the juice of the grape: the ferment also 
is natural to the grape, for it is developed from the albumen- 



196 LESSONS IN CHEMISTRY. 

like matter of the pulp. Since the alcoholic fermentation is a 
transformation of glucose, and no air is necessary for the change, 
this fermentation may continue in closed vessels ; in sparkling 
wines or champagnes part of the fermentation takes place in the 
bottle, and the carbon dioxide formed is dissolved in the liquid 
under pressure. All the carbon dioxide has escaped from still 
wines. The fermentation of apple-juice and the juices of other 
fruits, which yields cider and the various fruit-wines, is quite 
similar to the fermentation of grape-juice. Wines contain from 
seven to twenty per cent, of alcohol. 

Beer, ale, and porter are produced from grain, preferably from 
barley. Grain contains no glucose, but during the sprouting of 
the grain, a ferment, diastase, is formed, which subsequently 
converts the starch in the grain into a sugar, called maltose. 
The barley is moistened and kept at a temperature of about 15° 
until a sprout as long as the grain is formed. The sprouting is 
then arrested by heating the grain, which is now called malt, 
to about 50°, after which it is ground to a coarse powder and 
is ready for brewing. It is then cooked for several hours with 
water at a temperature of 60° : maltose is produced, and this as 
well as other nutritious matter of the malt is dissolved. The 
liquid thus formed is heated with hops to impart an aromatic 
flavor, and is then rapidly cooled ; after a little yeast is added, 
the wort is allowed to ferment at as low a temperature as pos- 
sible, until in a few days beer or ale is obtained, according to 
the proportions of substances used. 

Beer contains from two to five per cent, of alcohol, and ale a 
somewhat larger proportion, sometimes as high as ten per cent. 
As there is in this country no government inspection of malted 
liquors, beer is often adulterated by the substitution of various 
more or less injurious bitter substances for the hops, and of glu- 
cose for a part of the malt : glucose is not injurious, but it con- 
tains no nutritious matter, as is the case with malt. 

Spirituous liquors are not natural products ; they are dis- 
tilled from various fermented liquids, and are only dilute alcohol 
containing some flavoring matter. Brandy is distilled from wine ; 
whiskey from malted liquors of all kinds, derived from corn, rye, 



ALCOHOLS. 197 

oats, and even potatoes ; rum is distilled from fermented molasses 
from sugar-cane ; gin is dilute alcohol flavored with the essential 
oil of juniper-berries. These liquids contain from forty to sixty 
per cent, of alcohol. 



LESSON XXXVIII. 

ALCOHOLS (2). 

308. Propyl Alcohols, C 3 H 7 .OH. — A substance of this com- 
position exists in very small proportion among the products of the 
alcoholic fermentation. It is a liquid, boiling at 98°. When we 
examine the composition of the hydrocarbon propane, we will no- 
tice that the three carbon atoms are not similarly related : two are 
related to one other carbon atom, but the third is related to both 

of the first two. 

HHH 

H-C-C-C-H 

HHH 

Chemists have discovered two propyl alcohols, and indeed two 
modifications of every compound containing the radical propyl, 
C 3 H 7 . We account for this in our theory by considering that in one 
of these alcohols the hydroxyl group replaces one of the hydrogen 
atoms of either of the two carbon atoms which are related to only one 
other carbon atom, and we see that all these atoms of hydrogen are 
similarly situated. The alcohol so formed is normal propyl alcohol, 
that which we have briefly considered. In the other alcohol of 
the same composition the hydroxyl group is joined to that carbon 
atom which is related to two others : it is called isopropyl alcohol, 
and boils at 81°. We see, then, that there are two propyl radi- 
cals : propyl, CH 3 -CH 2 -CH 2 ; and isopropyl, CH(CH 3 ) 2 . Those 
alcohols in which the carbon atom which holds the hydroxyl group 
in the molecule is related to only one other atom of carbon, are 
called primary alcohols. Those in which the hydroxyl group is in 



198 LESSONS IN CHEMISTRY. 

relations with an atom of carbon which is related to two others, 
are called secondary alcohols. 

309. Butyl Alcohols, C 4 H 9 .OH. — Chemists have succeeded in 
preparing four different butyl alcohols. They are all liquids, with 
the exception of one, which is a crystalline solid, in whose mole- 
cule we believe that the hydroxyl group is held by a carbon atom 
which acts as the centre of a system ; around this atom are 
grouped three other atoms of carbon and their accompanying hy- 
drogen atoms. It is C(CH 3 ) 3 OH. It is called a tertiary alcohol, 
that being the name applied to those alcohols in which the carbon 
atom which brings the hydroxyl group into the molecule is related 
to three other carbon atoms. 

310. Amyl Alcohols, C 5 H n .OH. — There are now known seven 
alcohols of this composition. Two of them exist in the oily resi- 
due which is left in the distillation of brandy and whiskey, and 
they are therefore products of the alcoholic fermentation of glu- 
cose. This oily residue is called fusel oil : it has a peculiar and 
not altogether pleasant odor, and, besides some ordinary alcohol 
which it still retains, is a mixture of propyl, butyl, and two amyl 
alcohols. The first two may be isolated from the mixture by 
careful fractional distillation, but for the separation of the two 
amyl alcohols chemical means must be employed. The crude 
fusel oil is a valuable solvent for many substances, to dissolve 
which ordinary alcohol would otherwise be required. Butyl and 
amyl alcohols do not mix in all proportions with water, as do ethyl 
and propyl alcohols. 

We pour small quantities of methyl, ethyl, butyl alcohol from 
fusel oil, and amyl alcohol from fusel oil, into four separate plates, 
and light them. We find that the first is the most combustible, 
and that we can light the last only with difficulty. The flame of 
the methyl alcohol is almost colorless, but the brightness of the 
flames increases to the amyl alcohol, which burns with a bright 
light. The effect of an increased number of carbon atoms com- 
bined in the molecule is then to render the compound less volatile, 
and more difficult to inflame. 

311. Glycols. — We have learned that when ethylene gas, 



GLYCEROL. 199 

C 2 H*, is passed into bromine, a direct combination takes place, 
and ethylene bromide, C 2 H 4 Br 2 , is formed ; we have also acquired 
some idea of the molecule of this new compound. When ethylene 
bromide is boiled with a solution of potassium carbonate in water, 
carbon dioxide is given off, and, in addition to the potassium 
bromide which is formed, the liquid contains a new body. 

C 2 H*Br 2 + K 2 CO» + H-'O = C 2 H 4 (OH) 2 + 2KBr + CO 2 
This new compound, C 2 H 4 (OH) 2 , is formed by the replacement 
of the two bromine atoms of ethylene bromide by two hydroxyl 
groups. It is called ethylene alcohol, and after the potassium 
bromide has crystallized it can be separated by careful fractional 
distillation, and then forms a syrupy liquid having a sweet taste. 
Since it is a neutral body, and is a hydroxide, it is called an alco- 
hol, and it is a diatomic alcohol, because it contains two hydroxyl 
groups. It is the first member of a series of diatomic alcohols 
derived from the hydrocarbons of the series C n H 2n . From its 
sweet taste, Wurtz, its discoverer, gave it the name glycol, and the 
diatomic alcohols are often called glycols. 

312. Glycerol, or Glycerin, C 3 H 5 (OH) 3 .— The fats and fatty 
oils are complex compounds containing the radicals of certain 
carbon acids, called the fatty acids (§ 338), and the radical of the 
well-known substance glycerol. By the action of metallic hydrox- 
ides on these compounds, atoms of metal replace the glycerol 
radical, which combines with the hydroxyl groups which were be- 
fore in the metallic hydroxide molecules. At high temperatures, 
steam acts precisely like the metallic hydroxides, and glycerol 
is manufactured by distilling fats and oils in a current of super- 
heated steam ; that is, steam which has been passed through 
very hot pipes. The radical of glycerol having been found to 
be the group C 3 H 5 , we may represent the change thus : 

C 3 H5(fatty acid radical) 3 + 3HOH = C 3 H5(OH) 3 + 3H(fatty acid radical) 
Fat or Oil. Water. Glycerol. Fatty acid. 

Glycerol is a colorless, syrupy liquid, having a sweet taste. Its 
density is about 1.28. It freezes below 0°, and melts at about 
17°. When it is heated, it boils at about 280°, but is partially 
decomposed, producing a very irritating odor of a substance called 



200 LESSONS IN CHEMISTRY. 

acrolein. It may be distilled in a vacuum or in a current of 
steam. It dissolves in all proportions of water and alcohol. 

A molecule of glycerol contains three hydroxyl groups ; gly- 
cerol is then a triatomic alcohol. Each of these hydroxyl groups 
may be replaced by a monatomic atom or radical, and a large 
number of glycerol derivatives have been so formed. 

313. Nitroglycerin, C 3 H 5 (NO 3 ) 3 . — In a little beaker glass 
placed in ice- water, we have prepared a mixture of equal volumes 
of strong sulphuric and nitric acids ; into this cold mixture we 
pour a few drops of glycerol, and, after stirring for a few moments, 
we throw the contents of the beaker into another glass nearly 
filled with cold water. A few drops of an oily liquid separate, 
and fall to the bottom of the glass ; we pour off nearly all the 
water ; then, by means of a glass tube drawn out to a small open- 
ing, we remove a drop of the oil, and place it on the corner of a 
piece of paper ; when we light tbis, it burns with a bright flash. 
We now allow another small drop to fall on a nearly red-hot piece 
of sheet iron ; it explodes with a loud report. We place another 
small drop on an anvil, and when we strike it with a hammer 
there is another loud explosion. The oil is nitroglycerin ; it has 
been formed by the removal of three molecules of water from one 
molecule of glycerol and three molecules of nitric acid, and the 
union of the remaining groups, C 3 H 5 and 3N0 3 . 

C 3 H5(OH) 3 + 3HN0 3 = 3H 2 + C 3 H5(N0 3 ) 3 
When nitroglycerin explodes, carbon dioxide and water are formed. 
The energy of the explosion is due to the energy of motion of the 
atoms in a molecule of nitroglycerin being greater than that en- 
ergy in the carbon dioxide, water, and nitrogen, and during the 
explosion the excess of energy appears as heat and the energy 
necessary to convert the products into the gaseous state. Nitro- 
glycerin is used in blasting operations, but it is usually mixed with 
a very fine sandy earth, and the mixture is called dynamite. This 
is made into cartridges which are exploded by a fuse and deto- 
nating cap, Nitroglycerin is not a safe body to handle ; for experi- 
mental purposes we prepare only a few drops of it, and it is better 
to dry the compound by a few hours' exposure to dry, warm air 
before using it. 



SIMPLE ETHERS. 



201 



LESSON XXXIX. 



SIMPLE ETHERS. 



314. Oxides of Hydrocarbon Radicals.— In a glass flask we 
have cautiously mixed some methyl alcohol 
with about its own volume of strong sul- 
phuric acid; after adapting a cork and 
straight tube to the flask, we heat it, and 
soon a colorless gas is disengaged. We 
light the gas, and it burns with a rather 
bright flame. (Fig. 91.) It is methyl ox- 
ide ; the sulphuric acid has removed a mole- 
cule of water from two molecules of methyl 
alcohol, and the two methvl groups are held 
together by an atom of oxygen. 

2CH3.0H = (CH3)20 + H20 

Methyl alcohol. Methyl oxide. 

The density of methyl oxide compared to 
hydrogen is 23. The gas is converted into 
a liquid at a temperature of — 23°. It is 
soluble in water, alcohol, and ether. 

In this compound we are again shown that the methyl group 
acts like an atom of hydrogen. It is a monatomic radical, and the 
composition of methyl oxide is like that of water, excepting that, 
instead of two atoms joined by an atom of oxygen, two groups 
or systems are held to the oxygen atom. 

H-O-H H3C-0-CH3 

Water. Methyl oxide. 

The other monatomic radicals form similar oxides, and these 
oxides form part of a large class, called the simple ethers. 

315. Ethyl Oxide, (C 2 H 5 ) 2 0. — This compound, which is 
commonly called ether, may be formed by a number of reactions. 
In a strong glass tube we have sealed some ethyl iodide and silver 




Fig. 91. 



202 



LESSONS IN CHEMISTRY. 



oxide, and have heated the tube for several hours in boiling water. 
We now find that the black color of the silver oxide has changed 
to yellow, which is the color of silver iodide. A double decom- 
position has taken place, yielding silver iodide and ethyl oxide, 
which we may recognize by its odor and other properties when we 
cut open the tube. 

Ag20 + 2C 2 H^I = 2AgI + (C 2 H5)20 
Silver oxide. Ethyl iodide. Silver iodide. Ethyl oxide. 

In a rather large flask we have mixed some ninety per cent, 
alcohol with one and four-fifths times its weight of strong sulphuric 
acid. We adapt to this flask a cork having three holes ; through 
one passes a thermometer (t, Fig. 92) ; another gives passage to a 




Fig. 92. 



tube through which we may allow alcohol to flow from a reservoir, 
while to the third is fitted a delivery-tube connected with a con- 
denser through which flows a stream of ice-water. We now heat 
the flask until the thermometer shows that the liquid into which 



ETHYL OXIDE. 203 

the bulb dips has a temperature of 140° ; then we regulate the 
flame so that the temperature may not rise further, and start a 
small stream of alcohol from the reservoir. This alcohol is 
quickly changed to ether, which, together with the water formed, 
distils and collects in the receiving-bottle. 

The reaction which takes place in this operation is worth our 
study. When alcohol is mixed with strong sulphuric acid, water 
is formed, and an ethyl group is substituted for one atom of hy- 
drogen of sulphuric acid, producing a compound called ethyl sul- 
phuric acid. 

C 2 H5.0H + H 2 SO± = H 2 + (C 2 H^HSO± 
Alcohol. Ethylsulphuric acid. 

If this ethyl sulphuric acid is heated, it is converted into sul- 
phuric acid and ethylene gas, C 2 H* ; if it is boiled with water, it 
again yields alcohol and sulphuric acid ; but if it is boiled with 
an additional quantity of alcohol, the result is sulphuric acid and 

ether. 

(C 2 H5)HS0 4 + C 2 1R0H = H 2 SO + (C 2 H5) 2 
Ethylsulphuric acid. Alcohol. Ether. 

If methyl alcohol be used instead of ethyl alcohol, a mixed oxide 
of methyl and ethyl is produced. 

(C 2 H5)HSO± + CH».OH = H 2 SO + CHMD-C 2 !!* 

Ethyl oxide is a colorless, very mobile liquid, having a pleasant 
odor and a somewhat burning taste. Its density at 0° is 0.736 ; 
it boils at 34.5°. It will dissolve in nine times its weight of water, 
and one part of water will dissolve in thirty-six parts of ether. 
Ether dissolves small quantities of sulphur and phosphorus, and 
large proportions of bromine, iodine, fats, oils, and many other 
substances which are insoluble in water. 

Ether is very inflammable, and its vapor forms an explosive 
mixture with air. We suspend a heated coil of platinum wire 
over a little ether in a beaker, as we made a similar experiment with 
alcohol, and the slow combustion of the ether vapor develops so 
much heat that the platinum wire becomes hot enough to inflame 
the ether. 

The vapor of ether is very heavy, — 2.564 compared to air. We 



204 LESSONS IN CHEMISTRY. 

pour a little ether into a warm beaker, and then, holding the short 
end of a small siphon immediately above its surface, establish a cur- 

rent of ether vapor by drawing out the 

||| a * r by the mouth ; the heavy ether vapor 

|||[jl| 1 continues to flow through the siphon, as 

1 would a liquid, and, when we light it, will 

1 burn as long as any ether remains in the 

^HW \ beaker (Fig. 93). 

| l|[|jB jy| The inhalation of ether vapor produces 

I M l 0ijB^0 anaesthesia, and for this reason ether is 

h^ii'-'I'^iiii 1,1 ' 1 '^'"'- 1 llsBB^ 

illP 1 ^ largely employed as an anaesthetic in surgi- 

^ cal operations. 

Tig. 93. F 

The other monatomic hydrocarbon radicals form 

oxides corresponding to the oxides of methyl and ethyl : each of these con- 
tains two radicals related to one atom of oxygen. The diatomic hydrocarbons, 
ethylene and its homologues, form oxides containing one atom of oxygen and 
one molecule of the hydrocarbon, as in ethylene oxide, which may be obtained 
by heating ethylene bromide with silver oxide. 

? H ' Br + A g *0 = CH2 >0 + 2AgBr 
CH2Br CH2^ 5 

Ethylene bromide. Silver oxide. Ethylene oxide. Silver bromide. 

316. Chlorides, Bromides, etc. — The class of simple ethers 
includes the chlorides, bromides, iodides, sulphides, etc. of the 
radicals, and in general those compounds which correspond to the 
simple salts of the metals, — that is, those salts formed by an acid 
containing no oxygen. 

These compounds may be made by heating together the corre- 
sponding acid and alcohol. If methyl alcohol is heated with 
strong hydrochloric acid, a colorless gas, methyl chloride, CH 3 C1, 
is disengaged. 

317. Ethyl Iodide, C 2 H 5 I. — In a glass flask we put some 
alcohol with about two-thirds its weight of amorphous phospho- 
rus ; to this flask we fit a cork, through which passes a bottle- 
shaped tube called an adapter. We have partially closed the 
lower end of the adapter with some broken glass, and on this have 
placed a mixture of broken glass with a quantity of iodine equal 
to two and three-fourths times the weight of the alcohol. The 



ETHYL IODIDE AND BROMIDE. 



205 





upper end of the adapter is connected with a condenser so inclined 
that liquid may flow from it into the flask (Fig. 94). When we 

heat the flask, 
the alcohol 

boils, and its 
vapor, conden- 
sing, runs back 
through the iodine, which is dissolved and 

I brought gradually in contact with the amor- 

!■# phous phosphorus. Phosphorus tri-iodide, 

lm$ PI 3 > i s tnen f° rme d ; but this immediately 

reacts with the alcohol, forming ethyl iodide 
and phosphorous acid. The whole reaction 
is expressed in the equation 



Fig. 94. 



3C 2 H5.0H + P + I 3 = 3C 2 H5I + P(OH)3 
Alcohol. Ethyl iodide. Phosphorous acid. 

When the reaction has terminated, we agi- 
tate the contents of the flask with a dilute so- 
lution of sodium hydroxide, arid decant the 
aqueous liquid from the heavy oily layer of 
ethyl iodide. We then remove all traces of 
water from the latter by shaking it with 
some fragments of calcium chloride, and may 
further purify it by fractional distillation. 
Ethyl iodide is a colorless liquid, having 
at 0° a density of 1.975. It boils at 72°. 

318. Ethyl Bromide, C 2 H 5 Br, may be made by distilling a 
mixture of alcohol, potassium bromide, and sulphuric acid diluted 
with its weight of water, the substances being used in the pro- 
portions required by the equation 

C 2 H5.0H + KBr + H 2 SO* = C 2 H5Br + KHSO* + H 2 

It is a liquid having a pleasant odor, and boiling at 40°. It is 
sometimes used as an anaesthetic. 

319. By various means the hydrogen atoms of these simple ethers may be 
replaced by chlorine, bromine, or iodine, and compounds are then formed 
which we may consider as derived from other radicals. By the replacement 



206 LESSONS IN CHEMISTRY. 

of a hydrogen atom in ethyl bromide by a bromine atom, we may obtain 
either ethylene bromide, CH 2 Br-CH 2 Br, or an isomeric compound called 
ethylidene bromide, CH 3 -CHBr 2 . One of the more important of the sub- 
stances derived by the continued replacement of the hydrogen atoms of a 
simple ether, is chloroform. 

320. Chloroform, CHCP, may be made by passing methyl 
chloride mixed with chlorine over charcoal heated to about 200°. 
It is manufactured by distilling a mixture of alcohol or acetone 
with chlorinated lime, commonly called bleaching powder. The 
reaction is very complex. Chloroform and water condense to- 
gether in the receiver, and the choroform separates in a heavy oily 
layer, for it is hardly soluble in water. It is decanted, shaken 
first with water and then with a solution of potassium carbonate 
to remove impurities, and then distilled with calcium chloride, 
which removes the little water which it held in solution. 

Chloroform is a colorless liquid having a pleasant, stimulating 
odor, and a sweet, burning taste. Its density is 1.5, and it boils 
at 60.8. It is not inflammable, and communicates a green tint 
to a flame in which a drop of it is introduced on the end of a 
glass rod. It is used as an anaesthetic. 

321. There is a bromoform, CHBr 3 , a heavy, colorless liquid, and iodoform, 
CHI 3 , a yellow, crystalline solid. They are formed by the action of bromine 
and iodine on alcohol in presence of alkalies. Iodoform is used in surgery. 



LESSON XL. 

ALDEHYDES, CARBON ACIDS, AND KETONES. 

322. In a small beaker, we mix some ordinary alcohol with a 
little potassium dichromate and some strong sulphuric acid. The 
mixture becomes warm, and the red color of the potassium dichro- 
mate is changed to green. Potassium dichromate, a compound 
derived from chromic acid, contains much oxygen, and it is re- 
duced by the alcohol, which becomes oxidized. The peculiar odor, 
somewhat resembling that of apples, which is developed in the 



ALDEHYDE. — CHLORAL. 207 

beaker glass, is due principally to a substance called aldehyde. 
Its composition is C 2 H 4 0, and it represents alcohol in which the 
hydroxyl and one hydrogen atom are replaced by an atom of oxy- 
gen. 

CH3-CH2.0H + = CH3-CHO + H20 
Alcohol. Aldehyde. 

Aldehyde is made by distilling a mixture of alcohol, sulphuric 
acid, and potassium dichromate, and condensing the product in a 
receiver surrounded by ice. It is a very volatile liquid, boiling 
at 21°. 

323. To aldehyde correspond compounds of the same nature 
derived from each primary alcohol. In each of them the hydroxyl 
and one atom of hydrogen of the corresponding alcohol are re- 
placed by an atom of oxygen. 

324. There is an interesting derivative of aldehyde produced 
by the prolonged action of chlorine on absolute alcohol. It is an 
oily liquid, called chloral, and represents aldehyde in which three 
hydrogen atoms are replaced by three atoms of chlorine. 

CC1 3 -CH0 CH3-CHO 

Chloral. Aldehyde. 

Chloral, or trichloraldehyde, combines directly with water, form- 
ing a crystalline compound called chloral hydrate. This is much 
used in medicine for its sleep-producing properties. 

325. We may consider that an aldehyde is an oxidation product 
of an alcohol. If the oxidation proceed still further, the aldehyde 
is in its turn converted into an acid, and the conversion of an 
alcohol into an acid may take place without the previous forma- 
tion of an aldehyde. 

Indeed, a carbon acid may be considered as a primary alcohol 
in which the hydroxyl remains, but the two atoms of hydrogen 
related to the same carbon atom as the hydroxyl are replaced by 
an atom of oxygen. 

CH3-CH2.0H + O 2 = CH3-CO.OH + H20 
Alcohol. Acetic acid. 

We consider, then, that a carbon acid is a compound contain- 
ing the monatomic group of atoms CO. OH, and this group is 
often called carboxyl. As in all acids, the hydrogen of the hy- 



208 LESSONS IN CHEMISTRY. 

droxyl group may be replaced by metal, and salts are so formed. 
If there be only one carboxyl group in the acid, there can be only 
one series of salts ; but if there be two carboxyl groups, we can 
understand that either both or only one of the hydrogen atoms 
may be replaced, and there will be two series of salts, neutral 
salts and acid salts. 

326. Formic Acid, HCO.OH, is the acid formed by the oxi- 
dation of methyl alcohol. We have already seen that it may also 
be produced from hydrocyanic acid. It exists naturally in certain 
insects, and it takes its name from its existence in ants. It is 
made by distilling oxalic acid with glycerol, taking care that the 
temperature of the mixture does not rise above 100°. The oxalic 
acid then forms with the glycerol a compound which is again de- 
composed by the heat, and dilute formic acid distils, while the 
glycerol is regenerated. 

C 2 4H2 = HCO.OH + CO 2 

Oxalic acid. Formic acid. 

Formic acid is a colorless, very acid liquid, having a pungent 
odor. It freezes at 8.5°, and boils at 99°. It mixes with water 
in all proportions. To a little formic acid in a test-tube, we add 
some strong sulphuric acid, and gently heat the tube : an effer- 
vescence takes place, and we may light the escaping gas at the 
mouth of the tube. It is carbon monoxide, for the formic acid 
has been decomposed into that gas and water. 

HCO.OH = CO + H20 

By replacement of the hydrogen of the hydroxyl in formic acid, 
formates are produced : they are soluble in water, and yield carbon 
monoxide when heated with sulphuric acid. 

327. Acetic Acid, C 2 H 4 2 = CH 3 -CO.OH, is obtained in 
large quantities during the manufacture of charcoal by the distilla- 
tion of wood in closed vessels (§ 226). The liquids which con- 
dense in this operation consist of tarry matter, dilute acetic acid, 
wood-spirit, and some other substances. After the tar has been 
separated, the acid liquid is neutralized with lime, and a crude 
calcium acetate, generally called pyrolignite of lime, is so formed. 



ACETIC ACID. 



209 



This is mixed with sodium sulphate, and the sodium acetate and 

insoluble calcium sulphate formed are separated by filtration. 

Ca(C2H30 2 ) 2 + Na 2 SO± = CaSO* + 2jSTaC 2 H30 2 
Calcium acetate. Sodium acetate. 

The sodium acetate is then purified by crystallization, and, by 
heating it with strong sulphuric acid, is again converted into 
sodium sulphate and acetic acid which distils. 

328. Vinegar is a dilute acetic acid produced by the oxidation 
of alcohol. The oxidation is brought about by a minute organ- 
ized ferment which has the property of absorbing oxygen from the 
air and transferring it to the alcohol. The change is called the 
acetic fermentation : it does not take place in strong alcohol. In 
one method of manufacture, 
the dilute alcohol, or wine, is 
allowed to trickle over beech- 
wood shavings contained in 
a large cask having a double 
bottom and numerous perfo- 
rations for the circulation of 
air (Fig. 95). A large num- 
ber of these casks are placed 
in rows, and the shavings are 
first saturated with some beet- 
juice or sour wine in which 
the ferment is already devel- 
oped. The slow oxidation 
of the alcohol produces so 

much heat that the temper- Fig. 95. 

ature rises to 30°. It is 

usually necessary to allow the same liquid to pass twice through 
the cask before all of the alcohol is changed to acetic acid. 

329. Pure acetic acid is a corrosive liquid, having a pungent 
odor. Its density at 0° is 1.08 ; it freezes at 17°, and boils at 
118°. It is soluble in all proportions of water and alcohol. 

In a test-tube we neutralize a few drops of acetic acid with a 
fragment of solid potassium hydroxide : then we introduce a few 

14 




210 LESSONS IN CHEMISTRY. 

grains of arsenious oxide, and heat the tube. Dense white 
vapors, having a very unpleasant garlicky odor, are disengaged. 
These are due to the formation of a very poisonous compound 
called cacodyl. The test enables us to recognize an acetate. 

330. Acetates. — Acetic acid contains only one atom of hydrogen replaceable 
by metal, and the acetates must contain one atom of a metal united with one 
or more groups, C 2 H 3 2 , according to the atomicity of the metal. 

331. Sodium Acetate, NaC 2 H 3 2 + 3H 2 0, crystallizes in colorless prisms, 
which effloresce in dry air, and the water may be entirely driven out by heat. 
It is very soluble in water. 

332. Lead Acetate, Pb(C 2 H 3 2 ) 2 + 3H 2 0, is commonly called sugar of lead, 
owing to its sweet taste. It is made by dissolving lead oxide, PbO, in acetic 
acid. Solutions of lead acetate are capable of dissolving an excess of lead oxide, 
and when carbon dioxide is passed through the liquid, lead carbonate is precipi- 
tated, while the neutral acetate remains in solution. Lead acetate is poisonous. 

333. Copper Acetate, Cu(C 2 H 3 2 ) 2 + H 2 0, forms beautiful bluish-green 
crystals. Verdigris is a combination of copper acetate and cupric oxide, CuO. 

334. Acetone. — When acetates are strongly heated they are converted into 
carbonates, and a vapor is given off which may be condensed to a liquid. This 
has the composition C 3 H 6 0, and is called acetone. 

Ca(C 2 H 3 2 ) 2 = C 3 H60 + CaCO 3 

Acetone is found among the products of the dry distillation of wood. It is 
a colorless liquid of peculiar, not unpleasant odor. It boils at 56.5°, and mixes 
with water, alcohol, and ether in all proportions. It resembles aldehyde in many 
of its reactions. As it is formed when isopropyl alcohol, CH. 3 CH(OH).CH 3 , 
is oxidized, we write its formula CH 3 .CO.CH 3 . Analogous compounds, called 
ketones, are derived from other secondary alcohols. 

335. Before leaving acetic acid, we must study one interesting manner of its 
formation. We know that hydrochloric acid will convert hydrocyanic acid 
into formic acid ($ 262), If the hydrogen of hydrocyanic acid be replaced 
by a methyl group, CH 3 , methyl cyanide is obtained, CH 3 .CN : when this is 
treated in the same way, acetic acid results. 

CH 3 CN + 2H 2 = NH 3 + CH 3 COOH. 

The higher homologues of formic and acetic acids are obtained in the same 
way from the corresponding cyanides. When these are boiled with potassium 
hydroxide, the nitrogen atom is always changed for an atom of oxygen and 
the group OK, thus forming a salt of that carbon acid which contains one more 
carbon atom than the radical of the cyanide. 

336. The general formula of this series of acids is C n H 2n 2 . The higher 
members form part of the natural fats, and hence the series is generally called 
the series of fatty acids. The third acid is propionic acid, C 3 H 6 2 . There are 



ETHEREAL SALTS. 



211 



two butyric acids : one of them exists in butter, and the other may be obtained 
from isopropyl cyanide. There are three valeric acids, C 5 H 10 O 2 ; the most 
common exists in valerian root. It is a colorless liquid, having a strong and 
unpleasant odor. 



LESSON XLL 

ETHEREAL SALTS AND FATTY ACIDS. 

337. Ethereal Salts. — In a glass flask connected with a 
good condenser (Fig. 96) we distil a mixture of strong alcohol 




Fig. 96. 



with nearly twice its weight of strong sulphuric acid and three 
times its weight of crystallized sodium acetate. A colorless, 
volatile liquid, having a fragrant odor, condenses in the receiver. 
This body is ethyl acetate, and has been formed by the replace- 
ment of the sodium atom in sodium acetate by an ethyl group. 

NaC 2 H^0 2 + C 2 H*.OH + H 2 SO± = C 2 H5.C 2 H30 2 + NaHSO 4 + H20 
Sodium acetate. Ethyl acetate. 

Ethyl acetate is an ethereal salt or compound ether, and may 
be regarded as derived from acetic acid by the replacement of 
its basic hydrogen by the radical ethyl. Ethereal salts differ 



212 LESSONS IN CHEMISTRY. 

from metallic salts in that they contain hydrocarbon radicals 
instead of metallic atoms ; they resemble the latter in the man- 
ner in which they are formed, as well as in their ability to enter 
double decompositions. Thus r ethyl iodide and silver nitrate can 
be made to react and produce silver iodide and ethyl nitrate. 

C 2 H 5 I + AgNO 3 = Agl + C 2 H5.N0 3 
Ethyl iodide. Silver nitrate. Silver iodide. Ethyl nitrate. 

We gently heat some ethyl acetate with an alcoholic solution 
of potassium hydroxide ; the odor of the ethyl acetate disap- 
pears ; potassium acetate and alcohol have been formed. 
C 2 H5.C 2 H 3 2 + KOH = KC 2 H 3 2 + C 2 H5.0H 

338. Many of those ethereal salts in which both basic and 
acid radical are carbon compounds exist naturally in fruits, 
and impart to these their characteristic scents. Amyl acetate, 
C 5 H u .C 2 H 3 2 , for instance, has the odor of pears, and methyl 
butyrate and amyl valerate exist in pineapples and apples. 
These compounds may be prepared artificially by processes 
analogous to that described for ethyl acetate, or they may be 
made by passing hydrochloric acid gas through a mixture of 
the corresponding acid and alcohol. In this case water and a 
simple ether (chloride) are first formed, and the latter at once 
reacts with the carbon acid, the result being a compound ether, 
while hydrochloric acid is regenerated. 

339. Fatty Acids. — As the number of carbon atoms in the 
fatty acids increases, these substances are more oily in nature and 
less soluble in water. They are liquids at ordinary temperatures 
until the molecule contains nine atoms of carbon, C 9 H 18 2 ; the 
others are solids, and the melting point is higher as the com- 
position is more complex. Ethereal salts of these acids exist 
in various vegetable products and in animal secretions.. The 
peculiar odors of animals are due to fatty acid ethers. We must 
pass by the intermediate members of the series and study more 
particularly those which are most largely used in the arts. 

340. Palmitic Acid, C 16 H 32 2 , exists in palm oil, where the 
radical of the acid is combined with the radical C 3 H 5 of glycerol. 
It is manufactured by distilling palm oil in a current of super- 



FATS AND OILS. 213 

heated steam : glycerol and palmitic acid are formed, and the 
latter solidifies to a white mass on cooling. This mass is strongly 
pressed, to remove a liquid acid, oleic acid, which, existing also in 
a glycerol compound in the palm oil, is formed at the same time. 
The palmitic acid is then used for the manufacture of soap and 
candles. 

341. Stearic Acid, C 18 H 36 2 , forms a large proportion of tal- 
low, and may be made by decomposing that substance by super- 
heated steam. It is a white solid, fusible at 69°. It dissolves 
in alcohol and ether, and may be crystallized from its solutions. 
With the* exception of the alkaline stearates, the salts of stearic 
acid are insoluble in water. 

342. Oleic Acid, C 18 H 34 2 .— Olive oil contains the glycerol 
compound of an acid which does not belong to the series of fatty 
acids. It is an unsaturated carbon compound, and its molecule 
contains two atoms of hydrogen less than that of stearic acid. It 
is called oleic acid, and exists in many oils and fats, but always 
mixed with certain of the fatty acid compounds. It is an oily 
liquid, which freezes at 4°. 

343. Fats and Oils. — The natural fats and fatty oils are com- 
pound ethers in which a glycerol radical replaces the basic hydro- 
gen of the fatty acids. We have already seen that glycerol is a 
triatomic alcohol : it contains three hydroxyl groups, and in the 
fats and oils each hydroxyl group is replaced by a fatty acid less 
the hydrogen of its hydroxyl. The natural fats must, then, repre- 
sent three molecules of fatty acid and a molecule of glycerol. 
The names of these fatty bodies are derived from those of the 
fatty acids which take part in their formation. 

344. Palmitin, C 3 H 5 (C 16 H 31 2 ) 3 , may be extracted from palm 
oil which has been solidified by cold and then subjected to 
pressure to remove the liquid fatty matters. It is a white solid, 
melting at 60°. 

345. Stearin, C 3 H 5 (C 18 H 35 0' 2 ) 3 , is also solid ; it exists in the 
solid fats, such as tallow. 

346. Olein, C 3 H 5 (C 18 H 33 2 ) 3 , constitutes the greater portion 
of olive oil, almond oil, and other analogous oils. It is a liquid, 
which solidifies at 10°. 



214 LESSONS IN CHEMISTRY. 

Oils are usually classed as fat oils and drying oils. The first 
are such as do Dot solidify on exposure to air, but become rancid 
and acquire an unpleasant odor. They are numerous, and include 
olive oil, cotton-seed oil, oil of sweet almonds, peanut oil, and 
many others. The drying oils, of which the type is linseed oil, 
absorb oxygen and become thick and hard when exposed to the 
air ; they are used in the preparation of paints and varnishes. 

347. Saponification. — The decomposition of a compound 
ether by a metallic hydroxide, a decomposition which results in 
the formation of a metallic salt and an alcohol, is in chemical 
language called saponification ; however, a more restricted sense 
of the word implies the decomposition of a fatty body, with the 
formation of soap and glycerol. We boil some palm oil or olive 
oil with a solution of sodium hydroxide ; the oil disappears, and a 
soap has been formed, while glycerol is set free in the liquid. It 
is necessary that ordinary soaps shall be soluble in water, and we 
have already seen that the only ordinary metals which yield solu- 
ble salts with the fatty acids are potassium and sodium ; in other 
words, the alkaline metals (§ 341). Soap, then, is an alkaline salt 
of one of the higher fatty acids, generally palmitic and stearic, to 
which must be added oleic acid. Soft soaps are made with potas- 
sium hydroxide, while sodium hydroxide yields the hard soaps. 

In the manufacture of soap, the fat or oil is first boiled with 
a rather weak solution of sodium hydroxide, generally known as 
concentrated lye, and, when the mixture becomes pasty, enough 
strong caustic soda is added to saponify the fat completely. To 
separate the excess of water, common salt is added ; this dissolves 
in the water, causing the soap to come to the surface, for common 
soap is insoluble in salt water. The salty water, containing the ex- 
cess of alkaline hydroxide employed, is then drawn off, and the soap 
hardens on cooling. As it is not easy to separate from the waste 
liquid the glycerol formed in the reaction, it is more economical to 
decompose the fat by superheated steam, and boil with sodium hy- 
droxide the fatty acid which floats on the dilute glycerol. While 
soap is soluble in water, it is decomposed by a large quantity of that 
liauicj, a small quantity of alkaline hydroxide being set free, while 



SAPONIFICATION. 215 

the fatty acid becomes insoluble. The free alkali produces the 
cleansing effects, and the fatty acid forms the lather : we know 
that soap will not produce a lather if we use too little water. 
Ordinary soap is insoluble in salt water, but a soap which is sol- 
uble in salt water may be made from cocoanut oil ; it is called salt- 
water soap. It contains an alkaline laurate and myristate, lauric 
acid, C 12 H 24 2 , and myristic acid, C u H 28 2 , existing as glycerol 
ethers in cocoanut oil. 

348. Stearin candles are made from a mixture of solid 
fatty acids obtained by saponifying tallow by superheated steam 
and a small quantity of lime. The small quantity of insoluble 
calcium soap so formed is decomposed by sulphuric acid, and the 
oleic acid is separated from the solid acids by pressing the mass 
between warm plates. The oleic acid is used for the manufacture 
of soap. Certain fatty bodies, among them palm oil, are entirely 
decomposed by superheated steam, without the aid of lime. In 
this reaction the water acts as would either an acid or an alkaline 
hydrate, part of its molecule completing the basic molecule of 
glycerol, while the other part completes the acid molecule. 

C3H5(Ci6H3i02)3 + 3HOH = C3H5(OH) 3 + 3C 16 H3202 
Palrnitin. Glycerol. Palmitic acid. 

The saponification of fats and oils may be brought about by the 
action of strong acids, such as sulphuric acid, for the strong acid 
forms a new compound ether with the glycerol radical, and sets 
the fatty acid free ; the compound ether may then be again 
decomposed into glycerol and acid by the addition of water. 



216 LESSONS IN CHEMISTRY. 

LESSON XLII. 

CARBON ACIDS (3). 

349. Lactic Acid, C 3 H 6 3 , occurs in sour milk, and is pro- 
duced in the lactic fermentation of various sugars. It is usually 
made by allowing a solution of glucose to which some sour milk, 
a little old cheese, and some chalk have been added, to ferment 
in a warm place until the whole is converted into a solid mass 
of calcium lactate. This is purified by crystallization, and de- 
composed by the exact quantity of sulphuric acid required to 
precipitate the calcium as calcium sulphate. The solution is 
then separated by a filter, and evaporated on a water-bath. 
Lactic acid remains as a colorless, very sour, syrupy liquid, which 
is decomposed when heated. 

Lactic acid is propionic acid, C 3 H 6 2 , in which one atom of hydrogen is 
replaced by a hydroxyl group ; it is consequently at the same time an alcohol 
and an acid. Chemists have obtained another acid of the same composition, 
an isomeride of lactic acid, and the differences of the two are due to different 
positions of the hydroxyl group. The isomer is called hydr acrylic acid, because 
it is decomposed by heat into water and an acid called acrylic acid, C 3 H 4 2 . 

CH3-CH 2 -C0.0H CH3-CH(0H)-C0.0H CH2(OH)-CH 2 -CO.OH 

Propionic acid. Lactic acid. Hydracrylic acid. 

Lactic acid exists in three modifications which differ from one another in 
certain physical properties, but have identical molecular structure. (See 
Stereochemistry , p. 344.) 

350. Oxalic Acid, C 2 H 2 4 , exists naturally in many plants ; 
it gives the sour taste to sour grass, and at certain seasons is 
present in small quantities in rhubarb-leaves. It is a product of 
the oxidation of many vegetable matters : it may be made by 
boiling starch with rather dilute nitric acid, and evaporating the 
liquid. It is now manufactured by heating to 200° a pasty mix- 
ture of saw-dust and potassium hydroxide ; potassium oxalate is so 
formed, and is separated by treating the mass with hot water, in 
which it is quite soluble. The solution of potassium oxalate is 



OXALIC ACID. 217 

then mixed with milk of lime, which is calcium hydroxide, and 

insoluble calcium oxalate is formed, while the solution contains 

potassium hydroxide, which is used for another operation. 

K 2 C 2 0± + Ca(0H)2 = CaC 2 0* + 2K0H 

Potassium oxalate. Calcium hydroxide. Calcium oxalate. 

The calcium oxalate is decomposed by sulphuric acid, which 
forms insoluble calcium sulphate, and the solution of oxalic acid is 
evaporated until it is strong enough to crystallize. 

Oxalic acid forms large, colorless prisms, containing two mole- 
cules of water of crystallization. In dry air, these crystals efflo- 
resce, and the anhydrous acid may be obtained by carefully heating 
them to 100°. Oxalic acid dissolves in fifteen times its weight 
of cold water, and is also soluble in alcohol. When heated to about 
150°, it is decomposed with formation of carbon monoxide, carbon 
dioxide, formic acid, and water. 

2C 2 H 2 0* CO + 2C0 2 + CH20 2 + H 2 

Oxalic acid. Formic acid. 

We have already learned that both carbon monoxide and formic 
acid are prepared by the decomposition of oxalic acid. 

351. We neutralize a solution of oxalic acid by the addition of 

a little ammonia-water, and then pour into it some solution of 

calcium chloride. A white precipitate of insoluble calcium oxalate 

is formed. 

(NH*) 2 C 2 4 + CaCl 2 = 2NH*C1 -f CaC 2 0* 

Ammonium oxalate. Calcium chloride. Ammonium chloride. Calcium oxalate. 

Oxalic acid is poisonous ; its antidote is chalk, which is calcium 
carbonate : this causes the formation of insoluble calcium oxalate. 

We have prepared some silver oxalate by adding solution of 
silver nitrate to a solution of oxalic acid neutralized with ammonia. 
The insoluble silver oxalate is separated by filtration and dried. 
When we heat a small quantity of this powder in a test-tube, it 
suddenly explodes, being decomposed into carbon monoxide, carbon 
dioxide, and silver. 

Oxalic acid consists of two carboxyl groups, and is therefore a dibasic acid, 

CO.OH 

i With monatomic metals it may form two series of salts, acid oxalates, 

CO.OH. 

in which only one atom of hydrogen is replaced by metal, and neutral salts, in 



218 LESSONS IN CHEMISTRY. 

which both atoms are so replaced. One atom of a diatomic metal like calcium 
will of course replace both hydrogen atoms. 

With the exception of the oxalates of potassium, sodium, and ammonium, 
the neutral oxalates of the metals are insoluble in water, but they are decom- 
posed by dilute sulphuric and hydrochloric acids. 

352. Tartaric Acid, C 4 H 6 6 , is the acid of grapes. In the 
casks in which wine is kept there is deposited an impure potas- 
sium acid tartrate, called argot. This is purified by crystallization 
from boiling water, and the product so obtained constitutes cream 
of tartar. By boiling the aqueous solution with chalk, and add- 
ing sufficient calcium chloride to form potassium chloride with 
the potassium, insoluble calcium tartrate is formed, while carbon 
dioxide is given off, and potassium chloride remains in solution. 
The calcium tartrate is separated by filtration, and, after being 
washed with water, is decomposed with the theoretical quantity 
of dilute sulphuric acid. Calcium sulphate is precipitated, and 
when the filtered solution has been sufficiently concentrated by 
evaporation, crystals of tartaric acid are formed. 

Tartaric acid is in large, prismatic crystals, soluble in about 
half their weight of cold water, and also soluble in alcohol. By 
the action of heat it is converted into several other acids, of which 
the compositions depend on the temperature at which the tartaric 
acid is decomposed. 

353. We can easily understand the molecular constitution of tartaric acid 
by studying that of substances to which it is intimately related. When ethy- 
lene cyanide (CN)CH 2 -CH 2 (CN) is boiled with potassium hydroxide, ammonia 
is disengaged, and there is formed the potassium salt of succinic acid, so called 
because it is formed by the action of heat on amber. 

CH2-CN CH2-CO.OH + 2NH8 

CH2-CN + 4H2 ° = CH2-CO.OH 
Ethylene cyanide. Succinic acid. 

There exists in apples, gooseberries, and many other fruits an acid called 
malic acid, and this has also been prepared artificially in such a manner as to 
show that it represents succinic acid in which one atom of hydrogen is replaced 
by a hydroxyl group. The replacement of two hydrogen atoms of succinic 
acid by hydroxyl groups yields tartaric acid. 

CH2-C0.0H CH(OH)-CO.OH CH(OH)-CO.OH 

CH a -CO.OH CH 2 -CO.OH CH(OH)-CO.OH 

Succinic acid. Malic acid. Tartaric acid. 



CITRIC ACID. 219 

Tartaric acid is, then, a diatomic alcohol, for it contains two hydroxyl groups 
related to two carbon atoms, and it is a dibasic acid, for it contains two car- 
boxyl groups, CO. OH. There are two series of tartrates, acid tartrates, in which 
only one atom of basic hydrogen is replaced, and neutral tartrates, in which 
both are replaced. 

Like lactic acid, tartaric acid exists in several modifications differing in 
certain physical properties. (See Stereochemistry, p. 344.) 

354. Potassium Acid Tartrate, KC 4 H 5 6 , is cream of 
tartar, and is made by purifying argol. It is almost insoluble 
in cold water, but dissolves in boiling water. When heated to 
redness, it leaves a residue of charcoal and potassium carbonate, 
which may be dissolved from the mass by water. Pure potas- 
sium carbonate is usually obtained in this manner. 

355. Potassium Tartrate, K 2 C 4 H 4 6 , is made by adding 
potassium carbonate to a boiling solution of cream of tartar as 
long as carbon dioxide is disengaged. When the concentrated 
solution cools, the salt separates in crystals which are very soluble 
in water. 

356. Potassium Sodium Tartrate, KXaC*H 4 6 , is com- 
monly called Rochelh salt. It is made by neutralizing with so- 
dium carbonate a boiling solution of cream of tartar. It forms 
beautiful, colorless crystals, freely soluble in water. 

357. Potassium Antimonyl Tartrate, K(SbO)C*H 4 6 , 
known as tartar emetic, is formed when antimonous oxide is boiled 
with cream of tartar. Its crystals contain one molecule of water 
of crystallization for every two molecules of the salt, and effloresce 
in dry air. It is soluble in water, and is poisonous. When hy- 
drogen sulphide is passed through its solution, an orange-colored 
precipitate of antimony sulphide is formed. 

358. Citric Acid, C 6 H 8 7 , exists in lemons, oranges, currants, 
and many other fruits. It is made by allowing the juice of 
lemons or sour oranges to stand until it begins to ferment, and 
then neutralizing the boiling filtered liquid with chalk. The in- 
soluble calcium citrate formed is washed with boiling water, and 
decomposed by dilute sulphuric acid ; citric acid crystallizes from 
the solution separated from the insoluble calcium sulphate. 

Citric acid forms large colorless crystals, soluble in about three- 



220 LESSONS IN CHEMISTRY. 

fourths their weight of cold water, and having a very sour taste. 
Its cold solutions are not precipitated by lime-water, but become 
turbid when the liquid is boiled, for calcium citrate is more 
soluble in cold than in hot water. Magnesium citrate is em- 
ployed as a purgative in medicine. 

Citric acid is a tribasic acid and a monatomic alcohol : its molecular struc- 
ture is represented by the formula 

CH2.COOH 

C(OH).COOH 

CH2.COOH 



LESSON XLIII. 
CARBOHYDRATES. 



359. The principal constituents of plants are substances com- 
posed of carbon, hydrogen, and oxygen, the last two elements 
being present in exactly the proportions required for the forma- 
tion of water. For this reason they have been called carbo- 
hydrates, or hydrates of carbon. It should be remarked, how- 
ever, that certain compounds, recently discovered* contain 
different proportions of hydrogen and oxygen, although they 
present all the other characteristics of this class, and must be 
included in it. According to their compositions we may arrange 
the carbohydrates in three series, of which glucose, saccharose, 
and starch may be regarded as the types. 

Of the very numerous substances belonging to each of these 
series we can study but a few of the more common. 

360. Glucose, or Dextrose, C 6 H 12 6 , is the sugar of grapes, 
and constitutes the efflorescences seen on raisins and dried figs. 
It is generally associated with another sugar, fructose, having 
the same composition. Honey and the juice of many fruits owe 
their sweetness to a mixture of these sugars. 



* For example, rhamnose, which has the composition C 6 H 12 5 . 



FRUCTOSE OR LEVULOSE. 221 

Glucose is manufactured by boiling starch with a large quantity of water 
containing about one-half per cent, of sulphuric acid. The starch is not added 
until the liquid is boiling, and after about half an hour's cooking it is com- 
pletely converted into glucose. The sulphuric acid is then neutralized with 
chalk, and after the insoluble calcium sulphate has been separated by filtra- 
tion, the solution of glucose is concentrated until it will solidify to a crystalline 
mass on cooling. 

Glucose forms small, rounded, crystalline masses, which con- 
tain one molecule of water of crystallization for each molecule 
of glucose. AYhen cautiously heated, it melts, and again be- 
comes solid at 100°, all the water of crystallization being then 
expelled. Glucose dissolves in about its own weight of cold 
water, and the solution has a sweet taste. It is much employed 
in confectionery and syrups, but it is only about one-third as 
sweet as ordinary sugar. 

In a test-tube we boil a mixture of sodium hydroxide solution, 
potassium and sodium tartrate, and cupric sulphate : the result- 
ing deep-blue solution, known as Fclrfings solution, is not changed 
by heat, but when we add a little glucose to the boiling liquid, 
the color changes to yellowish red, and on standing red cuprous 
oxide is deposited. The glucose has reduced the cupric solution : 
glucose then acts as a reducing agent. To a solution of silver 
nitrate we add ammonia- water until the precipitate at first formed 
is just redissolved. Now on adding a little glucose and gently 
warming the tube a brilliant minor of silver is formed on its 
walls. In these reactions the glucose is oxidized and converted 
into complex acids. 

We already know that by fermentation glucose is decomposed 
into carbon dioxide and alcohol. 

361. Fructose, or Levulose, has the same composition as 
glucose, and resembles it in most of its properties. It is more 
soluble in water and alcohol than the latter, melts at 95°. and 
rotates the plane of polarized light to the left, while glucose 
turns it to the right. 

The glucose molecule contains an aldehyde group. CHO, and 
five hydroxyl groups ; that of fructose a carbonyl group, CO, 
and five hydroxyl groups. Both sugars are monosaccharides : 



222 LESSONS IN CHEMISTRY. 

the former is an aldehyde-alcohol, or aldose, the latter a ketone- 
alcohol, or ketose. They have recently been made by synthesis. 
362. Saccharose, or Cane-Sugar, C 12 H 22 O n .— This com- 
pound, which is ordinary sugar, is extracted principally from 
sugar-cane, sugar-maple, beet-root, and sorghum. 

Maple-sugar flows from incisions made in the bark of the maple. Sugar- 
cane, beet-root, or the plants from which sugar is to be extracted, are finely 
cut, and subjected to strong pressure, by which the juice is expressed. The 
liquid is then heated by steam in large boilers, and milk of lime (calcium 
hydroxide) is added to neutralize the natural acids of the juice and form in- 
soluble compounds with certain nitrogenized principles which are present. 
The dissolved lime is precipitated by a current of carbon dioxide. The syrup 
is then heated, and filtered through a layer of grained animal charcoal, and 
afterwards concentrated at as low a temperature as possible by boiling in large 
vessels in which a vacuum is made by pumps. When sufficiently concentrated, 
the syrup is run into cooling-pans, where it is continually stirred, so that the 
sugar may separate in small crystals, as granulated sugar. This is refined by 
being again dissolved and filtered through animal charcoal, after which the 
syrup must be again evaporated and crystallized. The granulated sugar is 
freed from syrup by rapid rotation in a cylinder of wire gauze, through which 
the syrup is thrown by centrifugal force. The still moist product, called soft 
sugar, is dried by being sifted on a revolving cylinder heated by steam and 
contained in a large, partially open drum through which a current of air is 
constantly passing. 

During the manufacture of sugar a part of that substance is by the action 
of the heat and water converted into glucose and fructose. 

C 12 H 22 O n + H 2 = C 6 H 12 6 + C 6 H 12 6 
Saccharose. Glucose. Fructose. 

The mixture of these substances can be crystallized only with great diffi- 
culty, and the uncrystallizable syrup constitutes molasses. The purest sugar 
is rock candy, and is obtained by stretching threads through a vessel con- 
taining a very concentrated syrup. The sugar then deposits in large crystals 
on the threads. 

Sugar is insoluble in ether and in absolute alcohol. Its crystals 
are anhydrous. It melts at 160°, and on cooling forms a hard, 
amorphous mass. At about 210° it is partially decomposed, 
yielding a brown, bitter substance known as caramel. It does 
not reduce Fehling's solution, but by long boiling is converted 
into glucose, which then effects the reduction. 

Solutions of cane sugar turn the plane of polarized light to 



STARCH. 223 

the right, and this fact is taken advantage of to estimate the 
amount of sugar in a solution. 

Upon warming with dilute acids, sugar solutions are inverted ; 
that is, they will rotate the plane of polarization to the left. 
This is due to the fact that the saccharose molecule is resolved 
into one of glucose and one of fructose. 

Cane sugar represents the d [saccharides, to which also belong 
its isomers maltose and lactose. 

363. Lactose is a hard, not very sweet substance which exists 
in the milk of animals, and is usually made by simply evaporating 
the whey left in the manufacture of cheese. It has the same 
composition as saccharose, but its crystals contain one molecule 
of water of crystallization to one of lactose, C 12 H 22 O u + H 2 0. 
It rotates the plane of polarization to the right. By boiling 
with dilute acids it yields glucose and its isomer, galactose. 

364. Starch is found everywhere in the vegetable kingdom, 
and constitutes the greater part of all grains, and of many tube- 
rous roots like the potato. It is obtained by reducing potatoes to 
a pulp, and washing this pulp in a sieve through which flows a 
stream of water. The fibrous matters, consisting of the torn cells 
of the potato, remain in the sieve, while the small particles of 
starch pass through and are deposited from the water, which is 
allowed to flow slowly down long inclined planes. From grains 
the starch is extracted by grinding the grain to flour, and knead- 
ing the flour in a sieve under running water. The starch passes 
through, as before, while the nitrogenized matter of the grain 
forms a soft, elastic mass, called gluten. 

The starch so obtained is simply separated from the vegetable 
cells in which it was formed. It occurs as a fine powder, in which 
microscopic examination reveals a peculiar granular structure. 
The size and shape of these granules vary with the source of the 
starch (Fig. 97) : they are from 2 to 185 thousandths of a millimetre 
in diameter. They are formed of concentric layers, and their 
structure becomes apparent when a little starch is dried at 100°, 
and, after moistening with a drop of water containing a trace of 
iodine, is examined by the aid of a microscope. The granules 



224 



LESSONS IN CHEMISTRY. 



then swell, and, as the exterior layers burst, the interior structure 
is exposed (Fig. 98). 

Starch is insoluble in water and alcohol ; but, when it is rubbed 
with water in a mortar with rough sides, a small quantity of the 





Fig. 97. Fig. 98. 

interior of the granules appears to dissolve. When it is boiled 
with a large quantity of water, the granules burst, and a turbid 
liquid is obtained on cooling ; this contains some soluble starch, 
and holds in suspension the insoluble starch. When heated with 
water to 60° or 70°, starch forms a gelatinous mass, called starch 
paste. We have already seen that starch develops a blue color 
with iodine ; and as starch is the test for iodine, so iodine is the 
test for starch. The blue color fades on heating. 

While the composition of starch is represented by the formula 
C 6 H 10 O 5 , there is no doubt that it must be multiplied by a large 
factor to express the molecular composition : we may write it 
therefore (C 6 H 10 O 5 ) n . 

Boiling with dilute acids converts starch into glucose. 

C 6 H ioo5 + H 2 = C6H1206 

Diastase, which is formed during the germination of grain 
(§ 307), converts starch into maltose. 

365. Dextrin. — When starch is heated to about 210°, it is 
changed into a body which is soluble in water, and which is not 
colored by iodine. It is a pale-yellow powder, called dextrin. 
Its solution is gummy, and is used as a mucilage. 

366. Gums. — These substances which are obtained from cer- 
tain plants are analogous in composition to saccharose. Gum- 



CELLULOSE. 225 

arable and gum-tragacanth contain compounds of this descrip- 
tion. 

367. Cellulose contains the same proportions of carbon, hydro- 
gen, and oxygen as starch. It is the matter which forms the 
walls of young cells in vegetables, and is deposited, together with 
other matters, in the older cells. Linen, cotton, paper, and the 
pith of certain plants are almost pure cellulose, which may be 
obtained by washing linen or cotton successively with dilute so- 
lution of potassium hydroxide, water, chlorine-water, acetic acid, 
alcohol, ether, and water. The insoluble matter left after these 
operations is cellulose. 

It is a translucent, white solid, having a density of about 1.3. 
It is not soluble in any of the ordinary solvents. It dissolves, 
however, in the blue liquid obtained by shaking copper with 
ammonia-water in contact with the air. By the action of strong 
sulphuric acid on cellulose, a gummy mass is obtained, which 
long boiling with water converts into fermentable glucose. 

When paper is soaked in a cold mixture of sulphuric acid with 
half its volume of water, and is then thoroughly washed and dried, 
it is converted into a semi-transparent substance, which is called 
vegetable parchment. This is the substance generally used for 
dialysis (§ 220). 

368. Gun-cotton is made by soaking cotton wool in a mixture 
of about equal volumes of strong nitric and sulphuric acids, and 
washing the product in running water until the last traces of acid 
are removed. After drying in the air, the substance has all the 
appearances of cotton, but is not as soft to the touch. It is very 
inflammable, and burns with a flash, leaving no residue. 

G-un-cottons, of which there are several varieties, are the nitric 
ethers of the cellulose radical. The most explosive variety is 
called pyroxylin, and is the hexanitrate, C 12 H u (N0 3 ) 6 4 ; it is 
insoluble in water, alcohol, and ether. A large volume of gas 
is produced by. the explosion, and attempts have been made to 
substitute this variety of gun-cotton for gunpowder. Another 
variety is produced when cellulose is treated for a short time 
only, with nitric and sulphuric acids. It is quite soluble in a 

15 



226 LESSONS IN CHEMISTRY. 

mixture of alcohol and ether, and the solution is employed, under 
the name collodion, in photography and in surgery. The soluble 
variety is a mixture of the tetra- and penta-nitrates. Celluloid, 
so much used as a substitute for bone and ivory, is made by in- 
corporating camphor with gun-cotton under strong pressure. 
Starch, dextrin, the gums, and cellulose are polysaccharides. 



LESSON XLIV. 
BENZENE DERIVATIVES (i). 

369. We have already seen that the unsaturated hydrocarbon 
benzene, C 6 H 6 , acts precisely like the saturated hydrocarbons, in 
that its compounds are formed by the replacement of its hydrogen 
atoms by other atoms or groups. We have learned that mono- 
chlorobenzene, C 6 H 5 C1, is formed in this manner by the replace- 
ment of an atom of hydrogen by one of chlorine. Since we may 
consider that the alcohols are formed by the replacement of one or 
more hydrogen atoms in the saturated hydrocarbons by the same 
number of hydroxyl groups, we can understand that a similar 
replacement in benzene should yield substances analogous to the 
alcohols. While these substances do resemble the alcohols in some 
of their chemical relations, they have at the same time certain 
other properties ; the hydrogen of their hydroxyl is more readily 
replaced by atoms of metal than is that of the alcohols. They 
are called phenols ; the most simple is that in which only one 
hydrogen atom is replaced by hydroxyl, and it is ordinary phenol, 
commonly called carbolic acid. 

370. Phenol, C 6 H 5 .OH. — This important compound can be 
prepared artificially from benzene, but it is always obtained from 
coal-tar, for it is one of the products of the destructive distillation 
of coal. After the benzene has been separated from the tar, that 
portion which distils during the fractional distillation between 
150° and 200° is collected separately, and is mixed with a satu- 
rated solution of sodium hydroxide. A compound in which the 



TRINITROPHENOL. 227 

hydrogen of the hydroxyl in phenol is replaced by sodium is so 
formed; this is dissolved in boiling water, and the solution sep- 
arated from the oily matters, which remain unaffected. The solu- 
tion of sodium phenate is then treated with hydrochloric acid, the 
reaction yielding phenol and sodium chloride. 

C 6 H 5 .ONa + HC1 = C 6 H 5 .OH + NaCI 

Sodium phenate. Phenol. 

The phenol is not very soluble in water, and when it has sep- 
arated is dried with calcium chloride, and distilled. The product 
is then cooled in a mixture of ice and salt, and the phenol forms 
crystals which are separated and drained. 

Phenol crystallizes in colorless needles, melting at 42° ; it boils 
at 183°. It has a characteristic odor, and a burning taste. It 
is only moderately soluble in water. Although it does not red- 
den blue litmus, it readily reacts with the metallic hydroxides, 
forming crystallizable compounds which in some respects resemble 
the salts. It acquires a more or less intense red color on exposure 
to air and light. Phenol is an exceedingly valuable agent for the 
destruction of low forms of life. It prevents putrefaction and 
decay of animal and vegetable matters, because it prevents the 
development of the minute germs of life which are the cause of 
such decompositions. Phenol is poisonous, and in a concentrated 
form is quite corrosive to living animal tissues. 

When bromine-water is added to even a very dilute solution 
of phenol, a yellow precipitate of tribromophenol, C 6 H 2 Br 3 (OH), 
is formed. A pine shaving dipped in phenol and then exposed to 
the air acquires a blue color. These properties aid us in identi- 
fying phenol. 

When two hydrogen atoms of benzene are replaced by hydroxyl groups, di- 
atomic phenols, usually called oxyphenols, result. They naturally have the 
composition C 6 H 4 (OH) 2 . Three oxyphenols are known, and we have already 
seen that we can understand these cases of isomerism by attributing to the 
hydroxyl groups different positions in the system of carbon atoms which are 
so intimately related together. 

371. Trinitrophenol, C 6 H 2 (N0 2 ) 3 .OH, commonly called picric 
acid, is obtained by boiling phenol with concentrated nitric acid. 
C6H5.0H + 3HN0 3 = C 6 H 2 (N0 2 ) 3 .OH + 3H 2 



228 LESSONS IN CHEMISTRY. 

It crystallizes from its solution in boiling water in small, lemon- 
yellow scales, which are not very soluble in cold water. It has 
an exceedingly bitter taste. It has acid properties, for the three 
nitro-groups, NO 2 , seem to make the hydrogen of the hydroxyl 
more readily replaceable by metal. Animal fibres, such as wool 
and silk, are dyed yellow in solutions of picric acid. 

372. Potassium Picrate, C 6 H 2 (N0 2 ) 3 .OK, may be made by adding potassium 
carbonate to a boiling solution of picric acid as long as carbon dioxide is dis- 
engaged. It forms long 3 T ellow needles, only slightly soluble in cold water. 

373. Ammonium Picrate, C 6 H 2 (N0 2 ) 3 .ONH 4 , is obtained in a similar manner 
by neutralizing picric acid with ammonia-water. It burns with a flash, with- 
out leaving a residue, and has been used in the manufacture of certain kinds 
of gunpowder and colored fires ($452). 

374. Nitrobenzene, C 6 H 5 .N0 2 . — When benzene is added in 
small portions to a cold mixture of strong nitric and sulphuric 
acids, and the liquid is constantly stirred, it dissolves, and when 
this solution is poured into cold water, a heavy, yellowish oil 
separates. This is nitrobenzene ; a hydrogen atom of benzene 
has been replaced by a nitro-group, NO 2 . 

C*6H6 + HNO 3 = C 6 H5.N0 2 + IFO 

Nitrobenzene boils at 205°, and has the specific gravity 1.2. 

It bas an odor resembling that of bitter almonds, and is used in 

perfumery, especially for imparting an odor to soap. It is 

manufactured in large quantities for the production of aniline. 

375. Aniline, C 6 H 5 .NH 2 . — If nitrobenzene be treated with a 
mixture capable of generating hydrogen, the nitro-group is re- 
duced, and converted into a group NH 2 . Almost all reducing 
agents produce this change, but in the arts iron and hydrochloric 
acid is used. The hydrogen eliminated from the acid by the 
iron, with formation of ferrous chloride, then reduces the nitro- 
benzene to aniline. 

C 6 H5.N0 2 + 3H 2 = C 6 H5.NH* + 2H 2 
Nitrobenzene. Aniline. 

The operation is conducted in large cast iron cylinders in which the nitro- 
benzene is continually stirred with the reducing mixture. The excess of acid 
is then neutralized with lime, and the aniline formed distilled while steam is 
passed through the liquid. 

Aniline is a colorless liquid, but becomes brown on long ex- 



ROSANILINE. 229 

posure to the air. It has an unpleasant odor, and an acrid, burn- 
ing taste. It boils at 184°. In water it is but slightly soluble, 
but mixes in all proportions with alcohol and ether. Although 
neutral to litmus, it combines directly with acids, forming crys- 
tallizable salts. 

Aniline represents ammonia in which one atom of hydrogen is replaced by 
the monatomic group C 6 H 5 . All of the hydrocarbon radicals are capable of 
replacing the hydrogen of ammonia, and the compounds so formed are called 
compound ammonias, or amines. Thus, methylamine, dimethylamine, and 
trimethylamine are formed respectively by the replacement of one, two, and 
three atoms of hydrogen in a molecule of ammonia. 

XH3 XH2.CH 3 NH(CH 3 ) 2 N(CH*)3 

Ammonia. Methylamine. Dimethylamine. Trimethylamine. 

The group C 6 H 5 , which is benzene less one atom of hydrogen, is called phenyl, 
and phenol is then phenyl hydroxide, while aniline is phenylamine. 

376. To a little aniline in a test-tube, we add a crystal of potas- 
sium nitrate, and then some strong sulphuric acid ; a bright-red 
color is produced. In another tube we mix some aniline with 
about twice its volume of strong sulphuric acid, and then drop in 
a small fragment of potassium dichromate ; a magnificent blue 
color is developed, and becomes violet when the mixture is diluted 
with water. A little bleaching-powder, that is, chlorinated lime, 
added to aniline produces also a violet color. Similar reactions 
are applied on a large scale in the manufacture of numerous 
coloring matters derived from aniline. 

377. Rosaniline. — The benzene of commerce is not pure, it 
contains much methylbenzene or toluene : when it is converted 
successively into nitrobenzene and aniline, a nitro-derivative of 
toluene is also formed, and this is reduced to methylaniline. just 
as the nitrobenzene is reduced to aniline. When such aniline is 
heated with oxidizing agents, both the aniline and the methyl- 
aniline lose hydrogen atoms, and the residues of the molecules 
combine, forming a complex body called rosaniline. 

C6H?N + 2C?H 9 X + 30 = C 20 H 21 X 3 O + 2H20 
Aniline. Methylaniline. Rosaniline. 

Large quantities of rosaniline are manufactured by heating 
commercial aniline either with arsenic acid or under pressure with 



230 LESSONS IN CHEMISTRY. 

nitrobenzene ; the oxygen of the arsenic acid or of the nitro- 
benzene, removes hydrogen from the aniline. 

Rosaniline is a colorless substance, but its salts have magnifi- 
cent colors and are used as dye-stuffs. The rich red coloring 
matter called magenta or fuchsine is a compound of one molecule 
of rosaniline with one of hydrochloric acid. If a hot saturated 
solution of this body be treated with sodium hydroxide, the color 
disappears ; sodium chloride is formed, and rosaniline separates 
as an almost colorless, crystalline precipitate. 

The hydrogen atoms of rosaniline may be replaced by various 
monatomic radicals, such as methyl, ethyl, phenyl. The com- 
pounds formed by three such replacements are more easily ob- 
tained than the others, and the salts of the resulting tri-substi- 
tuted rosanilines constitute a numerous, varied, and valuable 
class of coloring agents, known as the aniline dyes. 



LESSON XLV. 

BENZENE DERIVATIVES (2). 

378. The hydrocarbons derived from benzene by the replace- 
ment of its hydrogen atoms by groups such as methyl or ethyl, 
are capable of forming both phenols and alcohols ; for if the re- 
placement of a hydrogen atom by hydroxyl be in the benzene 
radical, a phenol would result, while an alcohol would be formed 
by such a replacement in the methyl or ethyl group. Methyl- 
benzene or toluene can thus form an alcohol, called benzyl alcohol, 
and three isomeric phenols, which are called cresols. 

C6H5-CH3 HO.C 6 H4.(CH 3 ) C6H5.CH2.0H 

Methyl-benzene. Cresols. Benzyl alcohol. 

Our time will permit the study of only a few of these com- 
pounds. 

379. Benzaldehyde, C 6 H 5 .CHO. — When chlorine gas is 



SALICYLIC ALDEHYDE. 



231 



passed through boiling toluene, benzyl chloride, C 6 H 5 -CH 2 C1. is 
formed, and by alkaline hydroxides this may be converted into 
benzyl alcohol, C 6 H 5 -CH 2 .OH. Just as ordinary alcohol may by 
slow oxidation be converted into aldehyde, benzyl alcohol is by 
the action of nitric acid converted into benzaldehyde. The latter 
body is interesting, because it is the essential part of oil of bitter 
almonds, so much used for flavoring. The oil of bitter almonds 
is, however, poisonous, for it contains hydrocyanic acid, which, to- 
gether with benzaldehyde and glucose, results from the action of 
water on a substance called amygdalitis existing in the almonds. 

380. Benzoic Acid, C 6 H 5 .CO.OH, exists naturally in gum ben- 
zoin, and is the product of the oxidation of benzyl aldehyde and 
benzyl alcohol. It may be easily prepared from gum benzoin, by 
gently heating some of that resin in a shallow dish, over which is 
pasted a piece of filter-paper. 

We cover the dish with a 
beaker, and the vapor of ben- 
zoic acid passes through the 
paper, on which and in the 
beaker it condenses in beauti- 
ful feathery tufts (Fig. 99). 

Benzoic acid crystallizes in 
colorless needles or thin plates. 
It melts at 121°, and boils at 
250°. It is not very soluble 
in cold water, but dissolves in 
about twelve times its weight 
of boiling water, and is also 
soluble in alcohol. It is an 
excellent antiseptic or pre- 
servative. 

381. Salicyl Aldehyde, C 6 H*(OH).CHO. — The pleasant 
odor of essential oil of meadow-sweet is due to a compound repre- 
senting benzaldehyde in which an atom of hydrogen in the ben- 
zene group is replaced by the radical hydroxyl. It is at the 
same time an aldehyde and a phenol. It is a colorless liquid. 




232 LESSONS IN CHEMISTRY. 

boiling at 196°. It is heavier than water. Oxidizing agents 
convert it into 

Salicylic Acid, C 6 H 4 (OH)CO.OH, in which the group CHO 
of the aldehyde has been changed to carboxyl, CO. OH. Salicylic 
acid is now manufactured by the action of carbon dioxide on 
phenol, or, more correctly, sodium phenate. We may represent 
the reaction 

C6H5.0H + CO 2 - C6H±(OH)CO.OH 
Phenol. Salicylic acid. 

Salicylic acid occurs as methyl salicylate in oil of wintergreen : 
the basic hydrogen, that of the hydroxyl group, is here replaced 
by a methyl group, CH 3 . 

C 6 H*(0H).C0.0H C 6 H4(OH).CO.OCH3 

Salicylic acid. Methyl salicylate. 

When oil of wintergreen is boiled with potassium hydroxide, 
potassium salicylate and methyl alcohol are formed. 

C6H*(0H)C0.0CH3 + KOH = C6H*(OH)CO.OK + CH 3 .OH 

Salicylic acid crystallizes in needles or prisms which are 
scarcely soluble in cold water, but very soluble in boiling water, 
alcohol, and ether. It is largely used as a preservative, and to 
some extent in medicine. 

382. Gallic Acid, C 6 H 2 (OH) 3 .CO.OH.— Salicylic acid repre- 
sents benzoic acid in which one atom of hydrogen is replaced by 
a hydroxyl group. It is a phenol and an acid. Gall-nuts, which 
are little excrescences produced by the sting of an insect on the 
leaves and twigs of certain species of oak, contain a substance 
which, by continued exposure to air and moisture, undergoes a 
sort of fermentation. When the liquid is pressed from the dark- 
colored mass, there remains a compound which may be crystallized 
from boiling water in long, colorless, silky needles. It is gallic 
acid, a compound which we may consider as benzoic acid in which 
three hydrogen atoms are replaced by three hydroxyl groups. It 
is colorless and odorless. When carefully heated, it is converted 
into a white volatile substance known as pyrogallol, or pyrog allic 
acid, while at the same time carbon dioxide is disengaged. 

C6H2(OH) 3 CO.OH = C 6 H 3 (OH) 3 + CO 2 
Gallic acid. Pyrogallol. 



TANNIN. 233 

Gallic acid is quite readily soluble in water. Its solutions, 
especially if an alkaline hydroxide be present, absorb oxygen 
from the air, and become dark in color. This last property is 
also common to pyrogallol, a solution of which is used as a re- 
ducing agent in photography. 

383. Tannin. — The well-known astringent properties of certain 
plants are due to the presence of compounds known as tannins or 
tannic acids, of which there appear to be a number of varieties. 
They possess the property of coagulating albumen and gelatin, and 
of forming black or nearly black precipitates with salts of iron. 
Tannin may be extracted from gall-nuts by placing the coarsely- 
powdered nuts in a funnel and pouring through them ether which 
is not free from water. As the ether runs through, it retains the 
coloring matters, while the water in the ether dissolves the tan- 
nin, and the aqueous solution separates from the layer of ether 
on standing. When this solution is evaporated at a gentle heat, 
the tannin remains as a light, very porous mass. It has an 
astringent taste, and is very soluble in water. When exposed to 
moist air or boiled with dilute sulphuric acid, it is converted into 
gallic acid ; a temperature of about 210° decomposes it, with for- 
mation of pyrogallol and carbon dioxide. These reactions indicate 
that tannin is related to gallic acid, and at least one of its varieties 
appears to be formed by the union of two molecules of gallic acid 
with the loss of one molecule of water. It is therefore digcdUc acid. 
2C 7 H 6 5 = H 2 + C 14 H 10 O 9 

Gallic acid. Digallic acid. 

The black mixture obtained by mixing solutions of tannin with 
ferric salts constitutes ink. A good black ink may be made by 
exhausting 100 grammes of powdered gall-nuts with 1.4 litres of 
water, and adding to the filtered liquid a solution of 50 grammes 
of gum arabic and 50 grammes of ferrous sulphate, each in the 
least quantity of water which will dissolve it. After stirring the 
mixture, it is allowed to stand exposed to the air until it becomes 
quite black. 

The operation of tanning, or the conversion of animal skins into 
leather, depends on the formation in the skin of an insoluble com- 



234 LESSONS IN CHEMISTRY. 

pound of tannin and the albuminoid matter of the skin. The 
tannin is derived from oak bark, which is ground to a coarse pow- 
der and piled in alternate layers with the skins in deep vats. 
The vats are then filled with water, and the skins are allowed to 
soak for a few weeks or months, until they have become thoroughly 
penetrated by the tannin. 

384. Camphors. — The highly aromatic solids that constitute 
the class of bodies called camphors are derived from paramethyl- 
isopropyl-benzene, or cymene ; it is benzene in which two atoms of 
hydrogen have been replaced, one by a methyl group, CH 3 , the 
other by an isopropyl group, C 3 H 7 . Its composition is therefore 

C 10 H H 

C6H6 CW(CH3)(C 3 H7) 

Benzene. Cymene. 

Cymene exists naturally in the essential oils of chamomile and 
thyme. It is a liquid having a pleasant odor. Its relations to 
the series of camphors are indicated in the following formulae : 

CiOHW Cymene. 

C 10 H 14 O, Thymol, or thyme camphor. 

C 10 H 16 O, Common, or Japan camphor. 

C 10 II 18 O, Borneol, or Borneo camphor. 

C l0 H 20 O, Menthol, or mint camphor. 

385. Thymol, C 10 H u O, is a phenol, one of the hydrogen atoms 
of the benzene group in cymene being replaced by the group OH. 
It exists in the essential oil of thyme, from which it may be ex- 
tracted in the form of large, colorless, crystalline plates, fusible 
at 51°. It has a pleasant but penetrating odor, and is an ex- 
cellent preservative, for it has antiseptic properties, destroying 
low forms of life. 

386. Common Camphor. C 10 H 16 O, sometimes called laurel 
camphor, because it is obtained from the camphor laurel, a tree 
which grows in China, Japan, and the island Formosa. It exists 
in all parts of the tree, but is extracted from the wood, which is 
chopped in small pieces and distilled with water. The camphor 
vapor condenses on rice-straw, with which the head of the still is 
filled : it is removed and purified by a new sublimation. It forms 
semi-transparent, crystalline masses having a strong, aromatic 
odor and a sharp, burning taste. Its density at 0° is 1. It melts at 



INDIGO. 235 

175°, and boils at 209° ; it is exceedingly volatile, and even at or- 
dinary temperatures it sublimes in the vessels in which it is kept, 
condensing in the upper part in brilliant, colorless crystals. It is 
almost insoluble in water, but dissolves readily in alcohol and in 
ether. When small fragments of camphor are thrown on the sur- 
face of clean water, they move around with curious gyratory move- 
ments, which are caused by the pressure of the camphor vapor 
given off in uuequal quantities from different parts of the surface 
of the fragments. The currents of vapor may be made evident 
by dusting a small quantity of the fine powder called lycopodium 
on the surface of the water. 

387. Borneol, C 10 H 18 O, is obtained from an aromatic tree of 
the Sunda Isles. It forms small colorless crystals, having an odor 
like that of camphor, but at the same time resembling that of 
pepper. It is insoluble in water, but dissolves in both alcohol and 
ether. Strong nitric acid converts it into ordinary camphor. 

388. Menthol, C 10 H 20 O, is the solid part of the essential oil 
of mint, in which it is mixed with a hydrocarbon having the same 
composition as oil of turpentine. It forms colorless crystals, 
fusible at 36°. 

389. Indigo, C 16 H 10 N 2 O 2 , is prepared from several indigo 
plants, which are cultivated principally in India. The leaves and 
stems of these plants are soaked in water for a day or two ; a sort 
of fermentation takes place, after which the liquid is expressed 
and agitated in contact with the air. A blue deposit forms ; it is 
collected and boiled with water in large copper vessels, and then 
drained, pressed, and broken up into the fragments in which in- 
digo occurs in commerce. Indigo results from the decomposition 
of one of its compounds which exists in the plant. 

The best indigo has a coppery appearance. It is not perfectly 
pure, but a small quantity may be purified by gently heating it in 
a small flask through which hydrogen is passed : the pure indigo, 
which is called indigotin, then sublimes and condenses in small 
crystals around the cooler portions of the flask. It is insoluble in 
water, alcohol, and ether, but dissolves in strong sulphuric acid, 
especially in fuming sulphuric acid. The dark-blue solution so 



236 LESSONS IN CHEMISTRY. 

obtained is commonly called sulphate of indigo, and is used in 
dyeing. 

390. White Indigo, C 16 H 12 N 2 2 .— When indigo is subjected 
to the action of reducing agents, such as sulphurous acid and hy- 
drogen sulphide, it is converted into a dirty-white substance called 
white indigo. If a mixture of indigo, ferrous sulphate, and milk 
of lime be shaken in a corked bottle, and allowed to stand for a 
day or two, an alkaline solution of white indigo is obtained, from 
which the latter may be precipitated by a current of hydrochloric 
acid gas. The white indigo is insoluble in water, but dissolves in 
alcohol and in solutions of the alkaline hydroxides. If a white 
cloth be dipped in the yellowish solution in the bottle, and then 
exposed to the air, it rapidly becomes blue. White indigo, which 
is a compound of hydrogen with indigo, is again converted into 
indigo on contact with the air, and the experiment with the cloth 
is an illustration of the manner in which in dyeing the insoluble 
blue indigo is deposited in the tissues of fabrics. 

Indigo has been obtained artificially by a number of interesting 
reactions, which will ere long permit the manufacture of this 
important dye-stuff from the hydrocarbons of coal-tar. 



LESSON XLVL 

NATURAL ALKALOIDS. 

391. The compound ammonias, derived from ammonia by re- 
placement of one or more of its hydrogen atoms by various groups 
or radicals, are powerful bases. They combine directly with acids, 
forming definite crystallizable salts. Thus, methyl ammonium 
chloride is as definite a body as ammonium chloride or potassium 
chloride. 

KC1 NH±.C1 NH3(CH3).C1 

Potassium chloride. Ammonium chloride. Methylammonium chloride. 

An immense number of compound ammonias have been formed, 



CONINE. 237 

and well studied, so that their molecular constitutions are perfectly 
known. Many plants contain principles most of which have not 
been obtained artificially, but which so much resemble the com- 
pound ammonias in their chemical relations that we believe them 
to belong to the same class. They all contain nitrogen, and it is 
to the nitrogen atom or atoms that are due the basic properties of 
the compounds which are called natural alkaloids ; that is, alkali- 
like bodies. Most of these substances are poisonous ; they all exert 
peculiar and active effects on the animal economy. 

392. The processes adopted for the separation of the alkaloids 
from the plants or vegetable products in which they occur, vary 
according to the solubility of the particular alkaloid and its salts 
in various solvent agents. The alkaloids do not occur in an un- 
combined state in the plants, but united with some natural acid 
with which they form salts. If the natural salt be soluble in water, 
an aqueous extract of the compound may be used for the prepara- 
tion of the alkaloid, but usually very dilute sulphuric acid is em- 
ployed ; sometimes the plant or product must be extracted with 
alcohol. The alkaloid is then . set free by the addition of milk 
of lime or other alkaline hydroxide, which will form a salt with the 
natural acid : sometimes the salt formed is insoluble in the liquid 
employed, while the alkaloid dissolves ; sometimes it is the alka- 
loid which is insoluble and the salt which remains in solution. 
These circumstances must be investigated, and such a process 
adopted as will allow the alkaloid to be entirely separated, and it 
can then be easily purified by crystallization. 

Two important natural alkaloids are liquid ; they are conine 
and nicotine. 

393. Conine, C 8 H 15 N, is the active principle of poisonous 
hemlock. It is extracted from the seeds, which are crushed and 
distilled with an alkaline hydrate. The conine distils, and is 
neutralized with sulphuric acid, which converts it into a sulphate, 
of which the solution is evaporated to a syrupy consistence, and 
then exhausted with a mixture of alcohol and ether. When the 
alcohol and ether have been evaporated, the conine sulphate is 
distilled with a strong solution of sodium hydroxide, and sodium 



238 LESSONS IN CHEMISTRY. 

sulphate is formed, while the conine set free condenses, together 
with a little water. It may be dried by calcium chloride. Co- 
nine is a colorless, oily liquid, having a disgusting odor. Only 
slightly soluble in cold water, and still less in hot water, it dis- 
solves freely in alcohol and ether. It is very poisonous. 

The molecular structure of conine has been determined : it is 
one of the few alkaloids that have been produced artificially. 

394. Nicotine, C 10 H U N 2 , exists in tobacco, probably in combi- 
nation with malic acid. It may be obtained by extracting tobacco 
with boiling water, evaporating the filtered solution until it becomes 
a pasty mass, and mixing this residue with about twice its volume 
of alcohol. The alcoholic liquid separates in two layers, of which 
the upper contains the nicotine : it is decanted, and the alcohol 
distilled off. From the residue the nicotine is set free by potas- 
sium hydroxide, and dissolved out by ether. The impure nicotine 
may then be converted into an oxalate by the addition of oxalic 
acid, and when this is decomposed by potassium hydroxide, 
tolerably pure nicotine is obtained. 

Nicotine is a colorless liquid, having an irritating and most 
penetrating odor. It is very soluble in water, alcohol, and 
ether. It boils at 241°. It is an energetic base, and is one 
of the most active poisons known. Tobacco contains from two 
to about seven per cent, of nicotine, the most esteemed varieties 
being those which contain the least. 

395. Theobromine, C 7 H 8 N 4 2 , is the alkaloid of cacao, and 
may be extracted from cacao beans. It is a white, crystalline 
powder, having a bitter taste, and is not very poisonous. 

396. Caffeine, C 8 H 10 N 4 O 2 , sometimes called theine, exists in 
coffee, tea, and several other plant products. It may be prepared 
by making a strong tincture of tea with cold alcohol, and precipi- 
tating the filtered liquid with basic lead acetate. The mixture is 
filtered, and freed from lead by a stream of hydrogen sulphide, 
after which it is again filtered, evaporated to a small volume, and 
while still hot is treated with potassium hydroxide. Caffeine then 
crystallizes out as the liquid cools. It forms long, brilliant, white 
needles, containing one molecule of water of crystallization, which 
is driven out by a temperature of 100°. It has a bitter taste ; 



MORPHINE. ATROPINE. 239 

it is not very soluble in cold water, but dissolves readily in hot 
water and in alcohol. 

Theobromine and caffeine are related to uric acid, C 5 H 4 3 X 4 , a body which 
is found in the excrements of birds and serpents, and in urinary calculi. By 
the replacement of two of its hydrogen atoms by methyl we obtain theobromine, 
while caffeine is the trimethyl derivative. 

397. Morphine, C 17 H 19 N0 3 .— Opium, which is the thickened 
juice of the unripe capsules of the opium poppy, contains several 
alkaloids, of which the most important is morphine. The natu- 
ral salts in which these alkaloids exist in opium are soluble in 
alcohol ; laudanum and paregoric are tinctures of opium. Mor- 
phine may be extracted by making a cold watery extract of 
finely-cut opium, evaporating the filtered liquid, and adding 
sodium carbonate to the still hot syrup. In the course of a 
day morphine deposits, and may be collected on a filter and 
dissolved in dilute acetic acid. The filtered solution is then 
decolorized by animal charcoal, and the morphine again precipi- 
tated by ammonia. Morphine is almost insoluble in water, and 
insoluble in ether. It is dissolved by hot alcohol, from which it 
separates in crystals containing one molecule of water. It has a 
very bitter taste. 

When nitric acid is added to a little morphine, an orange-red 
color is produced. Ferric chloride solution produces a blue color 
with morphine. Morphine forms easily-crystallizable salts ; the 
sulphate > chloride, and acetate are used in medicine ; these salts 
are soluble- in water. 

The principal alkaloids of opium, besides morphine, are codeine, 
C 18 H 21 N0 3 , and narcotine, C 22 H 23 N0 7 ; both are crystallizable 
solids. Codeine is morphine in which one hydrogen atom is 
replaced by methyl. 

398. Cocaine, 17 H 21 N0 4 , exists in coca leaves, which are much 
used as a tonic and stimulant in South America. The hydro- 
chloride is employed in medicine as a local anaesthetic. 

399. Atropine, or Daturine, C 17 H 23 N0 3 , is the alkaloid ob- 
tained from the deadly nightshade and the thorn-apple. When 
it is administered internally, or applied to the eye, it produces 



240 LESSONS IN CHEMISTRY. 

dilatation of the pupil, which continues until all of the alkaloid 
has passed from the system. It is exceedingly poisonous. 

400. Quinine, C 20 H 24 N 2 O 2 . — Cinchona bark contains several 
alkaloids, the more important of which are quinine and cin- 
chonine. 

These alkaloids are almost insoluble in water, and their sulphates are the 
forms in which they are principally employed in medicine. For the manu- 
facture of these salts, the bark is extracted with very dilute sulphuric acid, 
and milk of lime added to the clear solution to precipitate the alkaloids. The 
deposit, containing also calcium sulphate and lime, is collected, dried, and ex- 
hausted with boiling alcohol, which dissolves the alkaloids. When the filtered 
alcoholic solution is evaporated, the cinchonine crystallizes first, being least 
soluble, and the quinine is then neutralized with sulphuric acid and the solu- 
tion concentrated until the sulphate crystallizes. 

Quinine sulphate crystallizes in very bitter, delicate white 
needles, only slightly soluble in cold water, but dissolving in about 
thirty times their weight of boiling water. It dissolves readily in 
water containing a little free acid. When ammonia is added to 
its solution, the free alkaloid quinine is precipitated as a white 
powder, while ammonium sulphate is formed. Quinine is soluble 
in about its own weight of alcohol, and in twenty-two times its 
weight of ether. It is almost insoluble in water. 

401. Cinchonine, C 19 H 22 N 2 0, is deposited from the alcoholic 
solution in which quinine still remains in solution during its ex- 
traction from cinchona bark. Its properties much resemble those 
of quinine, from which, however, it may be distinguished by its 
insolubility in ether. Quinine sulphate supposed to contain cin- 
chonine sulphate is treated with a little ammonia-water, and then 
agitated with ether; any cinchonine present will remain undis- 
solved. 

402. Strychnine, C 21 H 22 N 2 2 .— The poisonous and medicinal 
properties of nux vomica are due principally to two alkaloids, 
strychnine and brucine. They are almost insoluble in water, and 
may be extracted from nux vomica by a process like that which 
serves for the separation of quinine. They are both exceedingly 
bitter, crystallizable solids, nearly insoluble in water. Strychnine 
is almost insoluble in alcohol and ether, but dissolves in chloro- 
form. Brucine is soluble in alcohol, and somewhat soluble in ether. 



METALS. 241 

If a small fragment of potassium dichromate is placed beside 
a crystal of strychnine, and both are touched with a drop of sul- 
phuric acid, a rich blue color is produced, which quickly changes 
to violet, purple, and red, and finally fades. 

Strychnine is a violent poison ; when taken even in compara- 
tively trifling quantities, it produces terrible convulsions, resem- 
bling those of tetanus. 



LESSON XLVII. 
METALS.— SPECTRUM ANALYSIS. 

403. The classification of the elements as metals and non- 
metals is more for the sake of convenience than for the indication 
of absolute properties of either class. We may, however, consider 
that certain general properties are peculiarly manifested by the 
metals : they are good conductors of heat and electricity ; they 
are capable of acquiring a brilliant lustre, which is called the 
metallic lustre. These properties are, however, more or less de- 
veloped in some of the elements which we have already studied. 
It is not so, however, with a chemical property : the metals are 
capable of replacing the hydrogen of the oxygen acids, forming 
salts. Some of these salts we have already studied, and we have 
seen how the combining power or worth of a metallic atom is 
indicated by the number of hydrogen atoms which it is able to 
replace in an acid. Yet even in this respect the metals and non- 
metals do not seem to be widely separated, for antimony, which is 
so closely related to phosphorus and arsenic by the compositions 
and chemical natures of its compounds, is also capable of forming 
a few salts. 

The physical properties of the metals are most varied. They 
are opaque ; but many of them can be reduced to sheets so thin 
that they allow the passage of a faint light whose color depends 
on the metal employed. Their densities vary from 0.59, that of 
lithium, to 22.4, of osmium ; their freezing points, from 39° below 

16 



242 LESSONS IN CHEMISTRY. 

0°, where mercury freezes, to about 2500°. Some, like manga- 
nese and chromium, are hard enough to scratch glass; others are 
soft enough to be scratched and even cut by the finger-nail, like 
potassium, sodium, and lead. Most of the metals are malleable and 
ductile; they can be beaten or rolled into sheets and drawn into 
wires. All the metals are insoluble in water. 

404. Natural State of the Metals. — The condition in which 
the metals are encountered in nature depends upon the other 
elements for which they have strong affinities. Some of them 
are often found in the metallic state : among these are gold, silver, 
copper, and bismuth ; they are then called native metals. 

In general, the elements which are more usually combined with 
the metals in their ores are oxygen, sulphur, and chlorine. Iron, 
zinc, and manganese are found as oxides ; iron, copper, lead, mer- 
cury, zinc, and silver, as sulphides; sodium and silver, as chlorides; 
calcium and magnesium, as carbonates ; aluminium, as silicate. 

The process adopted for the extraction of a metal must of course 
depend upon the nature of its ores : oxides are heated with char- 
coal ; carbonates are first heated to drive off carbon dioxide, and the 
resulting oxide is reduced by charcoal. Sulphides are roasted, — 
that is, heated in the air, — by which the sulphur is converted into 
sulphur dioxide, and passes off in that gas, while an oxide of the 
metal is formed. The methods employed for the reduction of the 
chlorides differ according to the metal. 

405. Alloys are the compounds or mixtures which the metals 
form with one another. In the molten state, many of the metals 
are capable of mixing with one another in all proportions ; but by 
certain precautions, and the use of the proper proportions of 
metals, many alloys become crystallizable, and assume the proper- 
ties of true chemical compounds : such compounds, of course, con- 
tain their respective metals in the proportions required by the 
atomic weights. The alloys of mercury are called amalgams. 

SPECTRUM ANALYSIS. 

406. When a beam of white light is passed through a prism, 
it is dispersed or separated into a spectrum consisting of all the 



SPECTRUM ANALYSIS. 243 

colors, from red to violet. If the beam be narrow and rectangular, 
such as is obtained by excluding all light except that which 
passes through a rectangular slit, and the spectrum be thrown on 
a white screen, the colors will not be confused, nor will they be 
distinctly separate, but will blend gradually from red to orange, 
yellow, green, blue, indigo, and violet. When any solid substance 
is heated to bright incandescence, it emits a white light, whose 
spectrum will contain all the prismatic colors. We have already 
had occasion to observe the dazzling whiteness of the combustion 
of magnesium and phosphorus. We have seen, also, that an alco- 
holic solution of boric acid burns with a green flame. We may 
prepare alcoholic solutions of sodium chloride, strontium chloride, 
and barium chloride, by shaking those substances in separate 
bottles with alcohol which is not too strong. When we burn 
these solutions, we find that the flame is colored yellow by the 
sodium salt, red by that of strontium, and green by that of 
barium, small quantities of the salts being carried into the vapor 
of alcohol, and volatilized by the high temperature of the flame. 
If on the end of a small platinum wire we introduce separately a 
little of each of these salts into the flame of an alcohol lamp, or, 
better, that of a Bunsen burner, we find that the same coloration 
is produced. If the light from such a flame be passed through a 
narrow slit, and then through a prism, we find that the color is 
invariable for each substance. The sodium salt produces not only 
a yellow light, but a particular shade of yellow, which, because it 
has passed through the straight slit, forms a peculiar spectrum, 
consisting of a single line of yellow light. If we keep the slit 
and prism in the same positions, it matters not what compound 
of sodium we introduce into the flame, the same yellow line is 
always produced, and on the same part of the screen. If we in- 
troduce a little lithium chloride into the flame, the latter will be 
colored red, and we will find on the screen a red line of a fixed 
and constant shade, and always in the same position. 

Analogous facts have been discovered for all the elements, and 
we may say, generally, that while the light emitted by an incan- 
descent solid depends upon the temperature, being first dull red, 



244 



LESSONS IN CHEMISTRY. 



then orange, yellow, and white, an incandescent gas or vapor, on 
the contrary, always emits light of a constant color, depending 
on the nature of the substance. Usually the spectrum of an 
element does not consist of a single line, one color only, but of 
several and sometimes many lines ; but in each case the spectrum 
is peculiar to the element. 

407. These principles have been applied in spectrum analysis 
for the detection of the elements, and the instrument employed 
is called a spectroscope. It consists of a narrow slit at the end 
of a metallic tube containing a lens (A, Fig. 100), by which 




Fig. 100. 



the rays of light are made to enter a prism in parallel lines : the 
light, liaving passed through the prism, is directed into a short 
telescope (B), by which the rays are again brought to a focus, so 
that the image may be examined by the eye. In order that the 
exact position of any line may be accurately observed and the 
line identified, the image of a small graduated scale illuminated 



SPECTRUM ANALYSIS. 245 

by a faint light (C) is reflected into the telescope from the side 
of the prism opposite to the slit. This scale corresponds to a 
similar graduated scale which we might make on a screen on 
which a spectrum is thrown. The substance of which the 
spectrum is to be examined is then heated on a platinum wire 
in a Bunsen-burner flame exactly opposite the slit. Sometimes 
electric sparks are passed between points of the substance, and if 
the latter be a gas it must be enclosed in a tube and rendered 
luminous by sparks passed through it from an induction coil. The 
spectra of most metals are very brilliant lines (see frontispiece), 
so brilliant that they entirely obscure the more faint and broader 
bands of the spectra of the non-metals ; for this reason, when we 
heat sodium chloride in the burner flame, we can only observe the 
spectrum of sodium by the spectroscope, although the chlorine 
must also produce its spectrum. 

Spectroscopic analysis is exceedingly delicate: 3,-o~ocr,Toir °^ a 
milligramme of sodium chloride introduced into the burner flame 
will cause the yellow sodium line to flash out for an instant. 
While studying spectra, several chemists have observed lines 
which were not produced by any substance then known, and have 
thus been led to the discovery of new elements, of which small 
quantities were present in the substances under examination. 

The study of the spectrum of the sun's light and the light of 
the stars has shown us perfectly the elements which exist in an 
incandescent state in the atmospheres of those far- distant bodies. 
Some lines in the spectrum of sunlight corresponded to those of 
no element known on the earth, and chemists concluded that this 
unknown substance existed in the sun's atmosphere, and named it 
helium* This same element has recently been discovered, by the 
aid of spectrum analysis, in a mineral called cleveite, another 
evidence of the distribution of the elements throughout the 
universe, and of the oneness in origin of all matter. 

* Helium is a colorless gas, having a density of 2. Its molecule, like that 
of argon, contains only one atom. Helium seems to have no affinity for other 
elements. It is the only gas which has not been liquefied. 



246 LESSONS IN CHEMISTRY. 

LESSON XL'VIIL 
METALLIC COMPOUNDS.— SPECIFIC HEAT. 

408. Before we undertake the study of the individual metals, 
we will pass in review some of the facts which we have already 
learned concerning metallic compounds, and will develop them by 
the consideration of new details. 

Oxides and Hydroxides. — All excepting a few of the metals 
combine directly with oxygen at various temperatures. Potassium 
is the only metal which is oxidized by cold dry air, and for the ox- 
idation of some metals a very high temperature is required. The 
number of atoms of oxygen and of metal which combine together 
depends on the atomicity of the metal. Two atoms of a mon- 
atomic metal combine with one atom of oxygen, while in the for- 
mation of a monoxide only one atom of a diatomic metal takes 
part. The oxides of lithium, sodium, and potassium are soluble 
in water, but in dissolving they form hydroxides, which we must 
admit contain a hydroxyl group. These hydroxides are sometimes 
called hydrates. 

K20 + H20 = 2KOH 

The hydroxides of these three metals are the alkaline hydrates, 
the metals being called the alkaline metals. Nearly all the oxides 
are capable, under certain conditions, of forming hydrates, con- 
taining one or more hydroxyl groups, and the oxidation or rusting 
of metals in moist air always results in the formation of hydroxides 
and not oxides. Calcium, strontium, and barium oxides (or hy- 
drates) are less soluble in water than those of lithium, sodium, 
and potassium. The other oxides are almost or entirely insoluble 
in water. 

Some metals form several compounds with oxygen, and those 
which contain the least oxygen are basic oxides, or bases ; they are 
capable of reacting with acids, forming water and salts in which 
the hydrogen of the acid is replaced by metal. Most of the oxides 



OXIDES AND HYDRATES. 247 

containing two atoms of oxygen also react with acids : the forma- 
tion of the salt is accompanied either by a disengagement of 
oxygen or chlorine, or the production of hydrogen dioxide. 

BaO 2 + 2HC1 = BaCl 2 + H 2 2 

MnO 2 + 2HC1 = MnCl 2 + 2H 2 + CI 2 

2Mn0 2 + 2H 2 S0 4 = 2MnSO* + 2H 2 + O 2 

The sesquioxides contain two atoms of metal and three of 
oxygen. They react with acids like the lower oxides, forming 
the corresponding salts and water. 

Fe 2 3 + 6HC1 = 2FeCl 3 + 3H 2 

A1 2 3 + 6HC1 = 2A1C1 3 + 3H 2 

Fe 2 3 + 3H 2 SO* = Fe 2 (SO*) 3 + 3H 2 

In some cases, however, these oxides act like acid radicals. 
Thus aluminium sesquioxide combines with sodium oxide, pro- 
ducing a salt-like body, aluminate of sodium. 

2A1 2 3 + 6XaOH = 2Al(ONa) 3 + 3H 2 

Ferric oxide combines with the more basic ferrous oxide to 
form the ferroso-ferric oxide, FeO.Fe 2 3 , which constitutes mag- 
netic iron ore. Spinel, which may be regarded as the type of 
these salt-like bodies, consists of the oxides of aluminium and 
magnesium, MgO.AFO 3 . 

The oxides containing one atom of metal combined with 
three or more atoms of oxygen correspond to metallic acids. 
They are capable of reacting with water or with basic oxides, 
forming well-marked acids and salts. Chromium trioxide, 
CrO 3 , corresponds to chromic acid, H 2 Cr0 4 = H 2 + CrO 3 ; 
manganese heptoxide, Mn 2 7 , to permanganic acid. HMnO*. 

When highly heated with charcoal or in a current of hydrogen, 
most of the metallic oxides are reduced to metal, while either 
carbon monoxide, carbon dioxide, or water is formed. The 
oxides of calcium, barium, strontium, magnesium, aluminium, 
potassium, sodium, and lithium are not reduced by hydrogen, 
and the first five are not reduced by carbon except with the aid 
of the electric arc. 



248 LESSONS IN CHEMISTRY. 

409. Sulphides. — Nearly all the metals combine directly with 
sulphur at certain temperatures, and the sulphides formed are an- 
alogous in composition to the oxides. The alkaline sulphides, and 
those of calcium, strontium, and barium, are soluble in water ; the 
others are insoluble. 

At temperatures depending upon the nature of the metal and 
the state of division of the sulphide, oxygen decomposes all the 
sulphides, sometimes forming sulphur dioxide and leaving a metal- 
lic oxide or even the free metal, sometimes oxidizing the sulphide 
to sulphate, according to the nature of the metal. If a mixture 
of potassium sulphate and powdered charcoal be heated to redness 
in a covered crucible, a porous black mass is obtained ; it contains 
potassium sulphide, and if it be broken up and thrown into the 
air, this sulphide is oxidized to potassium sulphate, producing a 
shower of sparks. 

K2S + 20 2 = K 2 S0 4 

Potassium sulphide. Potassium sulphate. 

410. Chlorides, Bromides, and Iodides. — With few excep- 
tions the metals combine directly with free chlorine ; since in its 
compounds with the metals, as in its compound with hydrogen, 
chlorine is a monatomic element, the number of chlorine atoms 
contained in a molecule of a metallic chloride is an indication of 
the atomicity of the metal. 

All the metallic chlorides are soluble in water, excepting silver 
chloride, mercurous chloride, and cuprous chloride: plumbic 
chloride is only slightly soluble. 

As a rule, the bromides are more soluble than the correspond- 
ing chlorides, and the iodides more soluble than the bromides. 

411. The color of the metallic compounds may be remem- 
bered by certain general principles. If both the corresponding 
oxide or hydroxide and the corresponding acid be colorless, the 
salts are also colorless. If either the acid or the oxide or hy- 
droxide be colored, the salts are colored. The salts formed by 
the same metal with colorless acids are of about the same color : 
with colorless oxides or hydroxides the same colored acid forms 
corresponding salts of about the same color. In many cases the 



SPECIFIC HEAT. 249 

color of metallic compounds depends on water of crystallization, 
as we have already seen (§ 55), and is lost when that water is 
expelled. 

SPECIFIC HEAT. 

412. The atomic weights of the metals cannot often be estimated from their 
vapor-densities, for many of them are volatile only at such high temperatures 
that it is impracticable, or even impossible, to determine the densities of their 
vapors. Some of the metals form volatile compounds with chlorine or with 
various hydrocarbon radicals ; and since the molecular weights of these com- 
pounds can be determined without difficulty from the densities of their vapors, 
we can arrive at the atomic weight of the corresponding meial. 

The compounds of the metals with oxygen, with chlorine, and with other 
bodies of course contain a fixed number of atoms of metal with a definite 
number of atoms of other elements of which the atomic weight is known. 
Thus, we know that for every sixteen parts of oxygen, potassium oxide con- 
tains 78.2 parts of potassium. We have already studied the reasoning by 
which we conclude that the atomic weight of oxygen is sixteen j how shall 
we determine whether the 78.2 parts of potassium represent one, two, or three 
atoms of that metal ? 

In order to raise the temperatures of equal weights of different substances 
through the same number of thermometric degrees, very different quantities 
of heat are required. If we expose one kilogramme of mercury and one kilo- 
gramme of water, both at 0°, to the same source of heat, we find that when 
the water will have been heated to 1° the mercury will be at 30°. If, on the 
other hand, we place one kilogramme of mercury at 100°, with some ice, in a 
vessel so constructed that all of the heat will be employed in melting the ice, 
we find that only one-thirtieth as much ice will be melted as if we put in the 
same vessel one kilogramme of water at 100°. The relative quantities of heat 
which are required to raise equal weights of different substances through the 
same number of thermometric degrees, are called the specific heats of the sub- 
stances. Water is the substance whose specific heat is chosen as unity, and 
the specific heat of any substance then represents the quantity of heat required 
to raise a given weight of the substance through one degree, compared with 
that which will raise the same weight of water through the same temperature. 
The specific heat of mercury is, then, ^ — 0.03333. On comparing the specific 
heats of the liquid or solid elements, it has been found that just in the same 
proportion that the atomic weight increases, the specific heat diminishes ; the 
specific heats are inversely as the atomic weights. The product of the specific 
heat of any liquid or solid element by its atomic weight should, then, always 
give the same figures. This important fact was discovered by Dulong and 
Petit, and is generally called Dulong and Petit's law: its import is evi- 
dently that the atoms of the different elements all possess the same specific 
heat. An examination of the figures expressing the quantities involved will 
show the facts on which the law is based : 



250 LESSONS IN CHEMISTRY. 

Name of Element. Atomic Weight. Specific Heat. Product. 

Lithium 7 0.9408 6.586 

Boron 11 0.5 5.5 

Carbon 12 0.46 5.52 

Sodium 23 0.2934 6.748 

Magnesium 24 0.2499 5.998 

Aluminium . 27 0.2143 5.786 

Phosphorus 31. 0.1887 5.850 

Sulphur 32 0.2026 6.483 

Potassium . . a 39.1 0.1695 6.500 

Zinc 65.2 0.0955 6.230 

Bromine . 80 0.0843 6.744 

Iodine ., . . 127 0.0541 6.873 

Mercury 200 0.0325 6.494 

The average of the products of the atomic weights by the specific heats is 
6.4 : however, while the product, which we may call the atomic heat, is always 
near the number 6.4, it varies within certain limits. Were it always 6.4, we 
could readily obtain 'the atomic weight of any element by dividing 6.4 by the 
specific heat; as it is, the figures expressing the specific heat enable us to 
choose between two numbers widely separated, and have in several cases indi- 
cated that the number which had been supposed to represent the atomic 
weight should be halved or doubled. 



LESSON XLIX. 
LITHIUM.— SODIUM.— POTASSIUM. 

413. Lithium, Li = 7. — The metal lithium is very widely 
diffused in nature, but is found only in small quantity. As a 
silicate or phosphate it forms an essential though minor con- 
stituent of a number of minerals, such as lepidolite and triphy- 
lite, and other compounds of it occur in spring waters and the 
ashes of certain plants. 

Metallic lithium is obtained by decomposing fused lithium chlo- 
ride, LiCl, a colorless soluble salt, by a current of electricity. It 
is a silver-white metal, and does not tarnish in dry air. Its den- 
sity is the lowest of any solid known, being about 0.58. It melts 
at 180°, and may be melted in contact with the air without be- 
coming oxidized : when heated to redness in the air or in oxygen, 
it burns with a dazzling white flame. Lithium soon becomes tar- 
nished in moist air, being converted into lithium hydroxide, 
LiOH : when it is thrown on the surface of water, the same 



SODIUM. 251 

hydrate is formed, the water being decomposed and hydrogen 
disengaged. The lithium salts are soluble in water, and are 
colorless unless the corresponding acid is colored. They com- 
municate a red color to the Bunsen-burner flame, and their spec- 
trum is characterized by a brilliant red and a more faint orange 
line (see frontispiece). 

414. Sodium, Na = 23. — Nearly forty per cent, of the im- 
mense quantities of the common salt which exists in the ocean, 
in deposits of rock-salt, and in brine springs, consists of sodium. 
The metal may be obtained by distilling a mixture of sodium 
carbonate and charcoal, as described under potassium (p. 254). 
Xa 2 C0 3 + 2C = 2Xa + 3CO 
A moie recent process consists in reducing sodium hydroxide 
by means of a mixture of iron and carbon which results when 
finely divided iron and gas tar are heated together. 

6XaOH + 2C = 2X a 2C0 3 + 3H2 + 2Xa 
Large quantities of sodium also are now manufactured by 
decomposing the hydroxide or chloride by means of the electric 
current. 

Sodium is a white metal, so soft that it can be cut and moulded 
like wax. It is lighter than water, its density being 0.97. It 
melts at 90.6°, and boils at a red heat. It can be melted in the 
air without taking fire. Its bright surface rapidly tarnishes in 
moist air, being converted into sodium hydroxide. It is preserved 
in bottles containing naphtha, by which it is 
protected from the air. When a small piece of 
sodium is thrown on water, chemical action at 
once begins ; the sodium melts and rushes about 
with a hissing noise. The reaction frequently 
terminates with an explosion by which small particles of sodium 
hydroxide are thrown out, and we must make the experiment at 
a safe distance from the eyes. If the motion of the sodium be 
arrested, the heat will accumulate sufficiently to ignite the escap- 
ing hydrogen. We float a piece of filter-paper on some water in 
a plate, and throw on this wet paper a small piece of sodium : it 
at once melts, and soon the hydrogen takes fire, burning with a 
flame tinged bright yellow by a little sodium vapor (Fig. 101). 




252 LESSONS IN CHEMISTRY. 

415. Sodium Hydroxide, NaOH, is the product of the reac- 
tion of sodium with water. It is manufactured by a number of 
processes : when sodium carbonate in rather dilute solution is boiled 
with milk of lime, sodium hydroxide passes into solution, while 
insoluble calcium carbonate is formed. 

Na2C0 3 + Ca(OH) 2 = 2Na0H + CaCO 3 
This operation is somewhat expensive, on account of the large 
quantity of water which must be boiled away from the sodium 
hydroxide. The Le Blanc process for the manufacture of sodium 
carbonate (§ 240) can with slight modifications be made to yield 
considerable quantities of an impure sodium hydroxide, which re- 
mains in solution after the sodium carbonate has crystallized. 

Much sodium hydroxide of an excellent quality is now manufactured from 
cryolite ($ 240). The powdered mineral is boiled with milk of lime, insoluble 
calcium fluoride and a solution of aluminate of sodium being obtained. 
AlE3.3NaF + 3Ca(OH) 2 = 3CaF2 + Al(ONa)* + 3H 2 
The filtered solution is then boiled with a new quantity of pulverized cryolite, 
and all the sodium is so converted into soluble sodium fluoride, while alumin- 
ium oxide is precipitated. 

Al(ONa)3 + AlF«.3NaF = AW + 6NaF 
When the precipitate has settled, the clear solution is drawn off and boiled 
with milk of lime : calcium fluoride is precipitated, while sodium hydroxide 
remains in solution. 

2NaF + Ca(OH) 2 = 2NaOH + CaF 2 

By whatever process it be obtained, the solution of sodium 
hydrate is evaporated to dryness, and subsequently fused in iron 
boilers out of contact with the air. It then forms a hard, white 
solid, which if left exposed to the air absorbs moisture and car- 
bon dioxide, becoming converted into sodium carbonate. It is 
very soluble in water, and very caustic. It is commonly known 
as caustic soda, and is employed in enormous quantities for the 
manufacture of soap. 

Sodium dioxide, Na 2 2 , is made by heating metallic sodium 
to about 300° in a mixture of nitrogen and oxygen gases in 
which the proportion of the latter is gradually increased. It is 
a yellowish solid which is decomposed by water into sodium 
hydroxide and oxygen. 

2Na 2 2 + 2H 2 = 4NaOH + O 2 



SODIUM CHLORIDE. 253 

Sodium dioxide is a powerful oxidizing agent, and is used for 
bleaching animal fibres. 

416. Sodium Chloride, NaCl. — This compound is common 
salt. It exists in numerous and immense deposits of rock-salt in 
many localities. It is found in salt wells and salt springs, and 
constitutes the greater portion of the solid matter of sea-water. 
The water of the Atlantic Ocean contains, according to the local- 
ity, from 32 to 38 grammes of solid matter per litre ; the water 
of the Pacific contains somewhat less, but the average proportion 
of common salt in each is about thirty grammes per litre. The 
other constituents of sea-water are principally chlorides and sul- 
phates of potassium, magnesium, and calcium, with small quanti- 
ties of bromides and iodides. When the water is evaporated, the 
sodium chloride separates first, while the other salts remain in 
more concentrated solution. In warm countries the evaporation 
is often accomplished by the heat of the sun and exposure to con- 
stant winds, in large shallow basins into which the water is either 
pumped or led by sluices from the sea. 

Sodium chloride crystallizes in cubes, which may be obtained 
of large dimensions and perfectly transparent, by the slow evapo- 
ration of a saturated solution. It is anhydrous, but the crystals, 
especially if small, usually retain in the spaces between them a 
small quantity of water, which is converted into steam and causes 
the crystals to decrepitate — that is, crack into small pieces — when 
they are heated. It is soluble in less than three times its weight 
of cold water, and in about two and a half times its weight of 
boiling water. It is insoluble in pure alcohol. It melts when 
heated to redness, and volatilizes at a higher temperature. 

417. Tests for Sodium. — Since all the ordinary sodium salts 
are soluble and colorless, none of the ordinary reagents pro- 
duce either precipitates or colors in their solutions. Hydrofluo- 
silicic acid yields a white precipitate of sodium silico-fluoride 
(§ 221). We may readily recognize the presence of sodium by 
the yellow color which its compounds communicate to the color- 
less flame of the Bunsen burner. 

418. Potassium, K = 39. — For a long time the principal 



254 



LESSSONS IN CHEMISTRY. 



source of potassium compounds was the potassium carbonate 
obtained from wood-ashes. Large quantities of potassium car- 
bonate are now obtained from the double chloride of potassium 
and magnesium, called, from its source, Stassfurt salt (§ 242). 
Metallic potassium is prepared by decomposing potassium car- 
bonate by carbon at a white heat. An intimate mixture of 
potash, lime, and carbon, obtained by heating crude argol* out 




Eig. 102. 

of contact with the air, is placed in an iron cylinder, and heated 
to whiteness. Carbon monoxide is disengaged, and the potas- 
sium vapor condenses in a flat receiver, from which the liquid 
metal runs into vessels containing naphtha (Fig. 102). Potas- 
sium occurs in commerce as round, brownish masses, kept 
under naphtha for the same reason that sodium is so preserved. 
It is quite soft, and yields readily to the pressure of the 
finger-nail. When freshly cut, it displays a brilliant surface, 
but this rapidly tarnishes by the action of the air. Its density 
is about 0.86 ; it melts at 62.5°, and boils at a red heat, emit- 
ting a green vapor. When heated in air it burns, forming 
the oxides K 2 and K 2 4 . In moist air it is converted into the 



Argol contains considerable quantities of calcium tartrate. 



POTASSIUM HYDRATE. 



ZDD 




hydroxide KOH. When a small piece of potassium is thrown 
in water, it decomposes the latter so violently that the hydrogen 
disengaged is at once ignited, and the potassium rushes about in 
the burning gas, whose flame is tinged violet by 
the metal (Fig. 103). The experiment termi- 
nates with a little explosion, for the globule of 
potassium hydrate formed is at a very high tem- 
perature, and when it cools sufficiently to come 
in contact with the water, there is a sudden for- 
mation of steam. 

419. Potassium Hydroxide, KOH, is prepared by boiling milk 
of lime with a rather dilute solution of potassium carbonate. As 
soon as the reaction has terminated, the solution of potassium hy- 
droxide is poured off the deposit of insoluble calcium carbonate, 
and is rapidly evaporated to dryness in iron or silver dishes. It is 
then fused, and cast in cylindrical moulds (Fig. 104), so that it 



Fig. 103. 




t t 



^r 



ml 



Fig. 104. 



usually occurs in commerce in round sticks. It commonly contains 
considerable quantities of lime, potassium carbonate, silicate, and 
other salts. It may be purified by dissolving it in alcohol in 
which only the hydroxide is soluble, decanting the clear solution, 
and fusing in a silver dish the residue from which the alcohol has 
been distilled. It is white and opaque, and has a destiny of 2.1. 
It melts at a red heat, and volatilizes at a higher temperature. It 



256 LESSONS IN CHEMISTRY. 

is exceedingly soluble in water, and, when exposed to the air, ab- 
sorbs moisture and carbon dioxide, deliquescing to a liquid con- 
sisting of a solution of the carbonate. It is very caustic and 
corrosive, rapidly destroying animal tissues. It is employed in 
making soft soap. 

420. Potassium Chloride, KC1. — This salt forms transpar- 
ent, colorless cubes, exactly resembling the crystals of sodium chlo- 
ride. It is found native in some localities, and, in combination 
with magnesium chloride, constitutes carnallite, KCl,MgCP -f- 
6H 2 (see § 242). It dissolves in about three times its weight 
of cold water, and in less than twice its weight of boiling water. 

421. Potassium Bromide, KBr, is employed extensively in 
medicine. It is usually made by adding to bromine enough 
strong solution of potassium hydroxide to almost decolorize the 
liquid. The reactiou yields a mixture of potassium bromide and 
potassium bromate. 

6KOH + 3Br2 = 5KBr + KBrO* + 3H20 
The mixture is evaporated to dryness, and then heated to red- 
ness, sometimes with the addition of a little powdered charcoal ; 
the bromate then loses its oxygen, and is converted into bromide. 
After cooling, the mass is dissolved in water, and the salt made to 
crystallize. Potassium bromide forms beautiful colorless cubes, 
having an intensely salty taste, and soluble in about one and a 
half times their weight of cold water. 

422. Potassium Iodide, KI, is prepared in exactly the same 
manner as the bromide, iodine being substituted for the bromine. 
It also crystallizes in colorless cubes having a salty and at the 
same time bitter taste. It dissolves in about two-thirds its 
weight of cold water, and the solution will dissolve large quanti- 
ties of iodine, becoming dark brown in color. Bath the bromide 
and iodide of potassium of commerce occur not in transparent 
but in white, opaque crystals : they contain a trace of free alkali. 
When the transparent crystals have been put in the market, they 
have found no sale, being supposed to be impure. 

423. Tests for Potassium. — Like the salts of sodium, most 
of the potassium salts are colorless and soluble, and their solutions 



SILVER. 257 

are neither precipitated nor colored by the ordinary reagents. 
Hydrofluosilicic acid produces a gelatinous white precipitate of 
silico-potassium fluoride. When the solution of a potassium salt 
is mixed with a strong solution of tartaric acid, a white crystal- 
line precipitate of cream of tartar soon separates. Platinic chloride, 
PtCl 4 , produces a yellow, crystalline precipitate of potassium 
chloroplatinate, (KCl) 2 PtCP. The potassium compounds impart 
a violet color to flame, but the color is rather delicate, and often 
masked by the presence of sodium or lithium : it is then examined 
through a blue glass which does not allow the passage of the 
light from the sodium and lithium flames, but through which the 
violet potassium flame is distinctly visible. 

424. Analogies of Lithium, Sodium, and Potassium. — When we compare 
together the compounds of the metals which we have just studied, we find 
that the three form a group presenting the most evident chemical analogies. 
They are monatomic metals, capable of replacing the hydrogen of acids, atom 
for atom. One atom of either metal will combine with one atom of chlorine, 
or with one hydroxyl group, but two atoms are required to combine with the 
diatomic atom of oxygen. Moreover, the corresponding salts of these metals 
are isomorphous : they crystallize either in exactly the same forms, or in forms 
which are easily derived one from the other. The rare metals caesium and 
rubidium form part of the group just considered. 



LESSON L. 
SILVER. Ag = 108. 

425. Silver is found in the metallic state, and in combination 
with many other elements, among the more ordinary of which are 
sulphur, chlorine, arsenic, and "antimony ; it is frequently asso- 
ciated with lead and copper. 

When the silver ores do not contain lead, the silver is extracted 
by amalgamating it with mercury and then driving off the latter 
by the action of heat. Several processes are employed ; in all of 
them the silver is first converted into silver chloride. The German 
method consists in roasting the powdered ore with common salt : 
the sulphides present are thus oxidized, while the silver is con- 
verted into chloride. The cold mass is pulverized, and washed 

17 



258 



LESSONS IN CHEMISTRY. 




Fig. 105. 



with water to remove all soluble salts formed ; the residue is then 
put into barrels with water and scrap iron, and these amalgamation 

barrels are rotated by machinery 
until the contents are thoroughly 
mixed (Fig. 105). Silver is set 
free, while the chlorine combines 
with the iron. Mercury is now 
introduced, and forms an amalgam 
with the silver. The liquid amal- 
gam is strongly pressed in canvas 
bags, and the greater part of the 
mercury is squeezed out. The 
semi-solid amalgam remaining is 
heated until the mercury is expelled, and the residue is metallic 
silver containing a certain proportion of copper derived from copper 
sulphide in the ore. This method is now obsolete. 

In the process adopted on the Pacific slope, the ore is reduced 
to a very fine powder, which is mixed with a proportion of com- 
mon salt depending on the amount of silver to be chloridized. 
By appropriate machinery, this mixture is thrown into a tall 
chimney-shaft through which a current of very hot air is rising. 
Under these circumstances, all the silver is at once converted into 
chloride, which falls to the bottom of the shaft, from which it is 
removed when about a ton has accumulated. It is then washed 
in a stream of water, and the insoluble silver chloride settles as 
a pulpy mass. This pulp is mixed with a little cupric sulphate 
and common salt in iron pans heated by steam, and about one 
hundred and fifty pounds of mercury are added for every ton of 
the pulp. After five or six hours' grinding, the mercury contains 
all the silver, which is reduced partly by the iron of the pan, 
partly by the conversion of some mercury into chloride. The 
amalgam is then agitated with water, and, after it is dried, the 
mercury is driven off by distillation in. cast-iron retorts. 

426. Galena, or lead sulphide, an important lead ore, often con- 
tains a considerable proportion of silver, which forms an alloy with 
the lead when the ore is reduced. Large quantities of silver are 






SILVER. 



259 



extracted from such lead by a process called, from the name of its 
inventor, Pattinsonizing. When .a melted alloy of lead and sil- 
ver containing even small quantities of the latter metal is allowed 
to cool, almost pure lead first solidifies in crystals : this is the fact 
on which the process is based. The molten lead is allowed to 
cool slowly, and, by means of large ladles, the crystals of lead are 
removed as fast as they are formed, so that the metal which 
remains liquid to the last is an alloy rich in silver (Fig. 106). 




Fig. 106. 



As the lead crystals so removed still contain a little silver, they 
are submitted a second and a third time to the same operation, so 
that pure lead is obtained on one hand, and a very rich silver alloy 
on the other. The lead is entirely removed from the alloy by 
a process called cupellation. The metal is melted on a shallow 
hearth swept by the flame of a small furnace. This hearth, which 



260 



LESSONS IN CHEMISTRY. 



is called a cupel, is covered by a sheet-iron dome (G, Fig. 107), 
which can be raised and lowered as necessary. When the whole 




Fig. 107. 

of the metal is melted, a blast of air is blown on its surface from 
pipes called tuyeres (t t), and the lead is oxidized. The oxide 
melts, and, being lighter than the metal, is drawn off through a 
notch cut in the side of the cupel, and the notch is gradually 
deepened as the level of the fused metal becomes lowered. The 
silver does not oxidize, and at last, when its surface is covered with 
only a thin layer of molten lead oxide, that layer breaks suddenly, 
and the brilliant surface of the silver appears with a flash. The 
blast of air is then stopped, and the silver is either drawn off into 
ingot-moulds or allowed to solidify in the cupel. 

427. Silver is the most brilliantly white metal. It is exceed- 
ingly malleable and ductile. Its density is 10.5. It does not 
tarnish on exposure to the air, but above its melting point, which 
is about 1000°, it absorbs or combines with about twenty-two times 
its volume of oxygen from the air. The oxygen is expelled vio- 
lently as the metal solidifies, and portions of the still liquid silver 
are often projected from the vessel, while its surface is thrown into 
curious tree-like forms. This phenomenon is called " spitting." 



SILVER CHLORIDE AND OXIDE. 261 

Ozone oxidizes silver to the dioxide Ag 2 2 . It is blackened by 
nydrogen sulphide, silver sulphide being formed on its surface ; 
the discoloration of silver-ware is due to traces of hydrogen sul- 
phide in the air ; the sulphur in eggs, mustard, etc., rapidly 
blackens silver spoons. Boiling sulphuric acid dissolves silver 
slowly, converting it into sulphate ; hydrochloric acid forms in- 
soluble silver chloride on its surface, and the metal beneath is so 
protected from further action. It dissolves readily in nitric acid, 
red vapors being disengaged and silver nitrate formed. It is not 
attacked by the alkaline hydroxides, and therefore silver vessels 
are used for the concentration and fusion of those compounds. 

428. Silver Chloride, AgCl, is one of the more important 
silver ores ; it is the mineral horn-silver, so called from its ap- 
pearance and somewhat elastic, horn-like structure. We have 
already seen that it is precipitated on the addition of hydrochloric 
acid or a soluble chloride to solution of silver nitrate. It then 
forms a white, curdy precipitate, which darkens and undergoes 
partial decomposition on exposure to light. If a piece of zinc be 
placed in some recently-precipitated and still moist silver chloride, 
the whole of the silver soon separates in the form of a gray 
powder, while zinc chloride is formed. Pure silver may be thus 
obtained, but for that purpose the silver chloride should be pre- 
viously well washed with dilute sulphuric acid, and the silver 
powder must be thoroughly washed by shaking it many times 
with water and then allowing it to settle. Pure silver may also 
be made by fusing the well-washed chloride with sodium carbon- 
ate ; carbon dioxide and oxygen are disengaged, sodium chloride 
is formed, and the silver remains as a button at the bottom of the 
crucible. When recently precipitated, silver chloride dissolves 
readily in ammonia-water, from which it is again deposited when 
the ammonia is neutralized by an acid. 

429. Silver Oxide, Ag 2 0, is made either by precipitating a 
solution of silver nitrate by potassium hydroxide, or by boiling 
well-washed silver chloride with potassium or sodium hydroxide so- 
lution. It is a brown powder, insoluble in water, and decomposed 
by heat into silver and oxygen. 



262 LESSONS IN CHEMISTRY. 

430. Silver Sulphide, Ag 2 S, is found native in small octa- 
hedral crystals. It is precipitated by the action of hydrogen sul- 
phide on solution of silver nitrate, and may be formed by the direct 
union of silver and sulphur at a slightly-elevated temperature. 

431. Tests for Silver. — In solutions of silver salts, hydro- 
chloric acid produces a white precipitate of silver chloride ; this 
precipitate is soluble in ammonia-water, and darkens in color when 
exposed to light. Potassium iodide solution gives a yellow pre- 
cipitate of silver iodide, Agl, which also darkens by the action of 
light, but is only slightly soluble in ammonia. Hydrogen sulphide 
precipitates black silver sulphide. Potassium chromate precipitates 
red silver chromate, Ag 2 CrO, in neutral solutions which are not 
too dilute. 

432. Silver-plating. — It is often desired to cover other metals or glass 
with a thin layer of silver. This may be accomplished in several manners. 
Copper objects may be silvered by rubbing them with a mixture of moist 
silver chloride and sodium carbonate, but the layer of silver so deposited is 
very thin. The metals are most readily and evenly silvered by connecting 
the object to be plated with the zinc pole of a voltaic battery and immersing 
it in a solution of silver and potassium double cyanide, made by boiling silver 
chloride in a solution of potassium cyanide. The positive pole of the battery 
is connected with a plate of silver, or silver coin, immersed in the same liquid. 
The silver solution then always retains its strength, for the metal dissolving 
from the positive electrode replaces that which is deposited on the article to 
be silvered. We may readily coat the interior of a test-tube with a thin layer 
of silver by pouring into it a solution of silver nitrate and sufficient ammonia- 
water to redissolve the precipitate first formed : we then add a few drops of a 
solution of tartaric acid, and place the tube in water heated to about 50°. A 
flat piece of glass may be silvered by the same liquid, which is then poured 
on in just sufficient quantity to cover evenly the perfectly-cleaned glass. The 
layer of silver so formed is very thin, and allows the passage of a violet light. 

433. Assaying of Silver. — The term assaying means determining the pro- 
portion of pure metal in either an alloy or an ore, but 
is now usually restricted to the first. Silver is alloyed 
with copper, and the alloy may be assayed either by a 
dry process — that is, one in which no liquid is employed 
— or by a wet process. The dry process consists in melt- 

FlG. 108. ing a small quantity of lead in a cupel, which is a little 

shallow cup made of compressed bone-ash and is very 
porous (Fig. 108). A weighed quantity of the silver coin or jewelry to be as- 
sayed is then wrapped in a small piece of paper and placed on the surface of 




SILVER ASSAY. 



263 



the melted lead, in which it is quickly dissolved. The cupel is heated in a 
muffle (A, Fig. 109) which fits into an opening in the side of a muffle-furnace. 
The muffle is open only at the exterior end, and has a slit in the arched top, 
so that the air is drawn through it by the draught of the furnace. The lead is 
oxidized by the air, and in presence of lead the copper of the alloy becomes 
also converted into oxide; the fused oxides are absorbed by the porous cupel, 
and as soon as their last traces disappear, the flashing of the bright silver 
surface indicates that the operation is finished. When cold, the button of 
pure silver is weighed. 

The wet assay is an example of volumetric analysis which we must study. 
We know that by the addition of a solution of common salt to one of silver 
nitrate, silver chloride is precipitated, and, since one molecule of sodium chlo- 
ride reacts with one molecule of silver nitrate, we find that 58.5 parts by weight 
of salt will precipitate exactly 108 parts of silver in the form of chloride. 

NaCl + AgXO 3 = AgCl - NaNO 8 
(23 + 35.5) (108 4- 14 -r 48) (108 + 35.5) (23 - 14 - 48) 

By carefully adding a solution of common salt to a solution of silver nitrate, 
we can tell when all the silver has been converted into chloride, for no more 
precipitate is then formed. Now, if we know 
how much salt we have added, we can easily 
calculate how much silver was pres- 
ent, because every 58.5 parts of salt 
used will represent 1 08 parts of silver 
precipitated. Let us make a solu- 
tion of salt of which each litre shall 
precipitate ten grammes of silver. 
Since 108 grammes of silver require 
58.5 grammes of salt, 10 grammes of 




silver will require 



58.5X10 

108 



= 5.417 



grammes of salt. We make such a 

solution, and we know that every 

cubic centimetre of it will precip- 

-. . 10 jjrainrnes , r 

itate 5 = 1 centigramme of 

1000 

silver. We now dissolve in nitric 
acid about a gramme of our alloy of 
silver, accurately weighed, and then 
- introduce our salt solution into a 
FlG. 109. burette (Fig. 110), which is a glass Fig. 110. 

tube having a stop-cock at the bot- 
tom, and graduated so that we may measure how much of the liquid we allow 
to run out. Then the salt solution is slowly dropped into the solution of silver 
nitrate, which is agitated so that the precipitate may quickly settle, until the 
instant arrives when a drop produces no precipitate. We then carefully read 



264 LESSONS IN CHEMISTRY. 

off the exact quantity of salt solution used, and calculate the amount of silver 
present in the quantity of alloy analyzed, each cubic centimetre of the salt 
solution representing 0.01 gramme of silver. 

The silver coins of the United States contain 90 per cent, of silver and 10 
per cent, of copper. 

434. Photography. — The chloride, bromide, and iodide of 
silver, being partially decomposed by the action of light, are em- 
ployed in photography. An image of the object to be photo- 
graphed being thrown on a glass plate coated with either of these 
sensitive salts, those portions on which the light falls are dark- 
ened, and metallic silver is formed ; the shades or dark parts of 
the image remain unaffected in proportion to the intensity of the 
shade : then when the plate is placed in a liquid capable of dis- 
solving the unaltered salts, a negative photograph is obtained ; that 
is, one in which the natural lights and shades are reversed. This 
negative being placed over a paper sensitized by some compound 
alterable by light, a positive picture is obtained, for the light acts 
through the transparent portions of the negative. We can easily 
make a sensitive paper by soaking a piece of soft white paper in 
a solution of common salt, and, after drying it, putting it in a 
solution of silver nitrate in a dark room. Silver chloride is thus 
formed in the paper. If now we have a negative or drawing on 
glass, we may make a photograph ; or we may copy some leaves 
by placing them on the paper, and, after pressing them down un- 
der a glass plate, expose the whole to the action of sunlight. In 
a quarter of an hour we remove the plate, and soak the paper in 
a solution of sodium thiosulphate (§ 106), which dissolves out the 
unaltered silver chloride : this is necessary, since the light would 
otherwise blacken the paper uniformly. After thoroughly wash- 
ing the paper in water, we have an exact copy of the negative or 
leaves employed. 



CALCIUM. 265 

LESSON LI. 

CALCIUM.— STRONTIUM.— BARIUM. 

435. These three elements form a group of metals of which the correspond- 
ing compounds not only present remarkable chemical analogies, but resemble 
one another in many physical properties. "We have already had occasion to 
notice, during the study of certain of their salts, that they are diatomic ele- 
ments, capable of replacing two atoms of hydrogen in the acids. 

The metals are obtained by decomposing their fused chlorides by a power- 
ful electric current. They are harder than lead, and their surfaces, which are 
brilliant when freshly filed, rapidly tarnish in moist air. They decompose 
cold water, forming hydrates while hydrogen is disengaged ; when heated in 
the air or in oxygen, they take fire and burn brilliantly. 

436. Calcium, Ca = 40, is the metallic radical of lime, marble, 
gypsum, etc. Its density is about 1.6. 

437. Calcium Chloride, CaCP, may be made by dissolving 
white marble in hydrochloric acid. It is now obtained in large 
quantities as an accessory product in the manufacture of sodium 
carbonate by the ammonia process. It crystallizes in large color- 
less prisms containing six molecules of water of crystallization. 
These crystals are deliquescent ; when they dissolve in water, in 
which they are very soluble, they produce a marked lowering of 
temperature. A mixture of equal weights of crystallized calcium 
chloride and snow or broken ice produces a temperature of — 45°. 
When heated, the crystals melt, and at 200° four molecules of 
water are driven out, but the other two are retained until the 
temperature reaches redness. As the anhydrous calcium chloride 
cools, it then solidifies to a hard, white, crystalline mass ; this is 
used for drying gases and liquids with which it undergoes no 
chemical reaction. Its solution in water develops considerable 
heat. 

A saturated solution of calcium chloride boils at 179.5°. The 
low cost of calcium chloride obtained in the ammonia-soda pro- 
cess has permitted the adoption of a new and very cheap process 
for the extraction of sulphur from the earthy matters with which 



266 



LESSONS IN CHEMISTRY. 



it occurs. The sulphur ore is immersed in a hot solution of cal- 
cium chloride of such strength that it boils at about 120° ; the 
sulphur then melts and runs out of the earthy matters, and may 
be drawn off as it collects below the hot liquid. 

438. Calcium Oxide, CaO. — This substance is universally 
known, and commonly called lime. It is manufactured by de- 
composing limestone, which is calcium carbonate, by the action 
of heat, but it is necessary that the products of combustion shall 
pass through the heated mineral, for calcium carbonate is decom- 
posed only at exceedingly high temperatures when heated in cov- 
ered vessels. Very primitive furnaces or lime-kilns are usually 
employed, resembling holes in the side of a hill : above an open- 
ing at the bottom a sort of grate is arranged, and on this the coal 
and limestone are thrown from the top. The fire is then lighted, 
and in about three days the kiln is burned out. A continuous 




Fig. 111. 

and more economical lime-kiln has an opening at the base for the 
removal of the lime, and about three metres above this opening 
there are others by which the flames from furnaces pass directly 
into the mass of limestone. As the lime is raked out at the bot- 
tom, the limestone descends, and more is thrown in at the top 
(Fig. 111). 



LIME. 267 

Lime occurs in hard, compact masses of a white or gray color : 
it is called quick-lime. It is infusible at the highest temperatures 
which we can produce by combustion, but may be melted, and 
even volatilized, in the electric furnace : the molten mass crys- 
tallizes upon cooling. When exposed to the air, it absorbs 
moisture and carbon dioxide, cracks, increases in volume, and 
crumbles to a white powder, which consists of a mixture of cal- 
cium hydroxide and calcium carbonate. When a mass of lime is 
sprinkled with water, the latter is absorbed ; in a short time the 
lime becomes so hot that steam is given off, and, if sufficient 
water be used, the whole falls to a bulky powder of calcium hy- 
droxide, Ca(OH) 2 , which is called slaked lime. Lime which de- 
velops much heat and increases greatly in volume by hydration 
is called fat lime, but if there be little heat produced, and the 
volume not greatly augmented, the lime is said to be poor lime ; 
it then contains considerable quantities of impurities. 

Milk of lime is calcium hydroxide, that is, slaked lime, sus- 
pended in water. If this white, creamy liquid be allowed to 
settle, the clear liquid obtained is lime-water. This is a solution 
of calcium hydroxide, which dissolves in about seven hundred 
times its weight of cold water. It is only about half as soluble 
in boiling water. When lime-water is heated, it becomes turbid 
from the separation of part of the hydroxide, which again dis- 
solves as the liquid cools. 

Large quantities of lime are employed in building operations. 
Ordinary mortar is a mixture of slaked lime and sand, the prin- 
cipal object of the latter being to prevent the shrinking of the 
mortar as it dries. Mortar hardens because the calcium hydroxide 
gradually absorbs carbon dioxide from the air, and the calcium 
carbonate formed, adhering strongly to the surfaces with which it 
is in contact, binds them together. It is possible that a small 
proportion of calcium silicate is also formed during the hard- 
ening. 

Cements, of which Portland cement * is an excellent type, are 



* Named from its resemblance to Portland stone. 



268 



LESSONS IN CHEMISTRY. 



made by calcining limestone with from ten to thirty per cent, of 
clay. Sometimes the clay exists naturally in the limestone; some- 
times it is added in the form of dried river-mud. Clay is a hy- 
drated aluminium silicate, and is rendered anhydrous by the action 
of heat. It is probable that at the same time a little calcium sili- 
cate and aluminate of calcium are formed. However that may be, 
the hard mass resulting from the calcination is pulverized, and 
the powder is cement, or hydraulic lime. When it is mixed with 
water, it sets, or hardens to a solid mass, in a very short time. It 
has the property of hardening under water, and is invaluable in 
submarine architecture. Its hardening is apparently due to the 
formation of a double silicate of aluminium and calcium. 

439. Chlorinated Lime, CaCl(ClO). — This compound, which 




Fig. 112. 



is intermediate between calcium chloride, CaCl 2 , and calcium hy- 
pochlorite, Ca(ClO) 2 , is manufactured on an extensive scale by 
passing chlorine gas over well-slaked lime placed in thin layers 
on shelves in masonry chambers (Fig. 112), care being taken that 
the temperature does not become too elevated. It is largely em- 



STRONTIUM. 269 

ployed as a bleaching and disinfecting agent, and owes this prop- 
erty to the facility with which it gives up its chlorine. It is 
decomposed by very dilute acids, even by the carbon dioxide of 
the air. 

CaCl(ClO) + CO 2 = CaCO 3 + ■ Cl» 

When thrown into water, it yields a solution containing calcium 
hypochlorite and calcium chloride. 

2CaCl(C10) = CaCl 2 + Ca(ClO) 2 
Chlorinated lime. Calcium hypochlorite. 

When it is heated, or when its solution is boiled, it is converted 
into calcium chloride and calcium chlorate. 

6CaCl(C10) = 5CaCl 2 -f Ca(C10 3 ) 2 

Chlorinated lime. Calcium chlorate. 

440. Calcium Carbide, CaC 2 , is a compound produced when 
a mixture of lime and carbon is heated in the electric furnace. 
It is a dark crystalline solid, having a density of 2.2. It reacts 
with water to form calcium hydroxide and acetylene (see p. 186). 

441. Tests for Calcium.— Solutions of calcium salts are not 
affected by hydrogen sulphide. In solutions which are not very 
dilute, sulphuric acid and the soluble sulphates produce a white 
precipitate of calcium sulphate. Solution of oxalic acid to which 
a few drops of ammonia have been added, yields a white precipitate 
of calcium oxalate, even in the most dilute calcium solutions. The 
salts of calcium communicate a reddish-yellow color to flame, and 
the calcium spectrum is quite characteristic. (See frontispiece.) 

442. Strontium, Sr = 87.5. — The principal strontium minerals 
are the sulphate, called celestite, on account of the blue color of 
many specimens, and the carbonate, called strontianite. The first, 
being the more abundant, serves for the preparation of the stron- 
tium salts : it is powdered, and intimately mixed with charcoal, 
and the mixture heated to bright redness in a covered crucible. 
Carbon monoxide is then disengaged, while the sulphate is re- 
duced to the sulphide, SrS. The gray mass containing this sul- 
phide is then treated with the acid corresponding to the desired 
salt, which separates in crystals when the solution is evaporated. 



270 LESSONS IN CHEMISTRY. 

443. Strontium Chloride, SrCP + ()H 2 0, crystallizes in 
deliquescent needles. It is moderately soluble in alcohol. 

444. Strontium Monoxide, SrO, is prepared by strongly 
calcining strontium nitrate. It is an infusible, gray, porous 
mass : when exposed to the air, it absorbs moisture and carbon 
dioxide. By the action of water, it is converted into strontium 
hydroxide, Sr(OH)' 2 , which is more soluble in water than calcium 
hydroxide, and crystallizes with eight molecules of water. There 
is also a dioxide, Sr0 2 . 

445. Tests for Strontium. — Solutions of the ordinary salts 
of strontium are colorless ; they are not precipitated by hydrogen 
sulphide. Sodium carbonate produces a voluminous white precip 
itate of strontium carbonate. Sulphuric acid precipitates stron 
tium sulphate in solutions which are not too dilute. Oxalic acid 
and ammonia produce a white precipitate of strontium oxalate. 
Flame is colored red by strontium compounds. 

446. Barium, Ba = 137. — Barium occurs in nature in heavy- 
spar, which is the sulphate, and witherite, which is the carbonate. 
The isolation of this metal is an extremely difficult matter : 
small amounts have been obtained by electrolyzing the fused 
chloride, and by reducing the oxide in the electric furnace. 
Its salts may be prepared by dissolving the native carbonate in 
the corresponding acid, or from the sulphate, which must first be 
reduced to sulphide. The finely-powdered sulphate is made into 
a paste with rosin and linseed oil, and the mixture is shaped into 
little balls which are calcined in a covered crucible. 

447. Barium Chloride, Bad 2 , is obtained when the sulphide 
is dissolved in hydrochloric acid, and the filtered solution suffi- 
ciently concentrated. Its crystals contain two molecules of water. 
They are soluble in rather more than twice their weight of cold 
water, in much less boiling water, and also slightly soluble in 
alcohol. Barium chloride is the reagent generally used for the 
detection of sulphuric acid. 

448. Barium Monoxide, BaO, often called baryta, is prepared, 
like strontium monoxide, by calcining the nitrate. It forms a gray, 
porous mass, which absorbs moisture and carbon dioxide from the 



BARIUM. 271 

air. If a fragment of this substance be sprinkled with a few 
drops of water, barium hydroxide is formed with such energy that 
the mass sometimes becomes red hot. 

449. Barium Hydroxide, Ba(OH) 2 , is made by dissolving the 
oxide in boiling water, which dissolves about one-tenth its weight. 
When the liquid cools, the greater part of the hydroxide is deposited 
in colorless crystals which contain Ba(OH) 2 -f- 8H 2 0. These crys- 
tals are soluble in water, and, under the name baryta-water, their 
solution is used for the precipitation of carbon dioxide as insoluble 
barium carbonate, or for the precipitation of sulphuric acid. 

450. Barium Dioxide, BaO 2 . — At a dull red heat, barium 
monoxide will absorb oxygen, and become converted into the di- 
oxide, which is made by passing oxygen over the monoxide heated 
in a porcelain tube or in a crucible. Barium dioxide is a grayish- 
white substance, which, when thrown into water, crumbles to a 
white hydroxide. It loses one atom of oxygen at a bright red heat, 
and the monoxide remains. By the action of strong sulphuric 
acid, barium dioxide is converted into barium sulphate, while ozone 
is disengaged. With hydrochloric acid, the hydrated dioxide yields 
barium chloride and hydrogen dioxide. 

451. Tests for Barium. — Hydrogen sulphide occasions no 
precipitate in solutions of barium salts. Sodium carbonate throws 
down white barium carbonate. Sulphuric acid precipitates insol- 
uble barium sulphate, even in exceedingly dilute solutions, and 
the precipitate is insoluble in nitric acid, either cold or boiling. 
Barium salts communicate a green color to flames. 

The barium salts are very poisonous. 

452. The nitrates of barium and strontium are employed in pyrotechny, for 
they impart to fireworks the characteristic flame colors of the metals. 

A red fire may be made by mixing 30 parts of potassium chlorate, 17 parts 
of sulphur, 2 of charcoal, and 45 of strontium nitrate. The materials must 
be pulverized separately, and may be mixed by repeated passing through a 
sieve. A green fire may be made by similarly mixing 33 parts of potassium 
chlorate, 10 of sulphur, 5 of charcoal, and 52 of barium nitrate. If it be 
desired that the fires shall produce little or no smoke, the following formulae 
may be used ; the ammonium picrate may be made by adding ammonia- 
water to a concentrated alcoholic solution of picric acid, until the liquid has 
an ammoniacal odor, and then collecting and carefully drying the precipitate. 



272 LESSONS IN CHEMISTRY. 

Ammonium Ferrous Strontium Barium 

picrate. picrate. nitrate. nitrate. 

Yellow . . o ... 50 50 

Green 48 ... ... 52 

Red 54 ... 46 

The stars for rockets and Roman candles are made by moistening the colored 
fires and forming them into small balls ; these are dried and introduced into 
the tube, from which they are projected by a small charge of gunpowder. 



LESSON LII. 
LEAD. Pb = 207. 



453. In many of its chemical relations, lead resembles calcium, strontium, 
and barium, and it might be classed in the same group of metals ; but in a 
number of its compounds it acts as a tetratomic element. It forms a dioxide, 
PbO 2 , and a tetrachloride, PbCl 4 . In the dioxides of strontium and barium, 
it is not probable that the atoms of these metals are tetratomic : it appears 
rather that the two atoms of oxygen are related to each other, while each is 
also related to the atom of metal. In lead dioxide the lead atom is tetratomic. 

454. The principal lead ores are galena, which is lead sulphide, 
and cerusite, which is the carbonate. The reduction of the latter 
mineral is an exceedingly simple process : it is heated with char- 
coal ; the reduced lead collects on the hearth of the furnace, and 
is drawn off as it accumulates. 

Galena may be reduced by heating it with scrap iron: iron 
sulphide and lead are formed, and, the lead being the heavier, the 
iron sulphide floats on the surface and is drawn off as slag. The 
more usual process, known as the reaction process, consists in 
heating the galena on the hearth of a reverberatory furnace (Fig. 
112) provided with openings (D) for the admission of air. Part 
of the lead sulphide is so converted into oxide, and another portion 
into sulphate. When this reaction has sufficiently advanced, the 
openings of the furnace are closed, and the heat is increased. 
Under these circumstances the unaltered sulphide reacts with both 
oxide and sulphate, metallic lead being formed, while sulphur 
dioxide is disengaged. 

PbS + 2PbO = 3Pb + SO 2 
PbS + PbSO* = 2Pb + 2S0 3 



LEAD. 273 

Sometimes charcoal powder is added after the air-openings are 
closed, in order to aid in the reduction of the oxide and sulphate. 




Fig. 112. 

Lead is a bluish-white metal, having a brilliant lustre, which 
soon tarnishes by exposure to air. It is soft, and can be scratched 
by the finger-nail : it is quite malleable, but has so little tenacity 
that it cannot readily be drawn into wire. Its density is about 
11.36 : it melts at about 334°, It may be crystallized by allow- 
ing a crucible full of the molten metal to cool until a crust forms 
on its surface, piercing the crust, and pouring out the still liquid 
interior. The interior of the crucible is then found to be lined 
with octahedral crystals. Molten lead absorbs oxygen from the 
air, and its surface becomes covered with a film of lead oxide, PbO. 

Lead is only slightly attacked by hydrochloric acid, and is 
scarcely affected by dilute sulphuric acid. Strong sulphuric 
acid dissolves it by the aid of heat, sulphur dioxide being given 
off. Nitric acid converts it into lead nitrate, and disengages 
red vapors. As the nitrate is almost insoluble in nitric acid, 
the latter should be diluted with water. 

Pure water containing dissolved air and carbon dioxide dissolves 
a small quantity of lead in the form of hydrate and carbonate, and 
for this reason lead is an unsafe metal for lining rain-water cisterns 
intended for storing drinking-water. Most spring- and river-waters 

18 



Zi4 LESSONS IN CHEMISTRY. 

contain small quantities of sulphates : when such water flows 
through lead pipes, the surface of the metal becomes quickly cov- 
ered with a film of insoluble lead sulphate, which protects the 
pipe from further action, and the water from being poisoned by 
the introduction of lead compounds. 

Lead, and all its soluble compounds, as well as such as may be rendered 
soluble by the juices of the stomach, are poisonous, and the poisonous effects 
are cumulative. Workmen employed in the manufacture of white lead, red 
lead, and other lead compounds, frequently suffer from chronic lead-poisoning, 
as do also painters and color-grinders. Small quantities of lead are then accu- 
mulated in the system, and cause peculiar disorders, among which lead colic 
is the most common : one of the characteristic symptoms of lead-poisoning is 
a peculiar blue line around the borders of the gums. The workmen in lead- 
works usually drink small quantities of an exceedingly dilute sulphuric acid, 
by which the lead in the system is converted into the insoluble and innocuous 
sulphate. In cases of chronic lead-poisoning, the administration of potassium 
iodide removes the metal from the tissues by the formation of lead iodide, 
which is soluble in solutions of potassium iodide, and can consequently be 
eliminated by the excretory organs. 

Metallic lead is used in the form of sheets for roofing and lining 
tanks ; it is manufactured into lead pipe ; type metal, which is 80 
per cent, lead and 20 per cent, antimony ; pewter, which contains 
between eighty and ninety per cent, tin, the remainder being lead; 
and plumbers' solder, an alloy of lead and tin. Enormous quantities 
of lead are employed for the manufacture of shot, which is made 
by allowing the molten metal to run through a sieve, and the drops 
to fall from a height into water. In common qualities of tin plate, 
a large proportion of the coating is lead instead of pure tin. 

455. Lead Chloride, PbCP, is prepared by boiling lead 
oxide in hydrochloric acid, and is precipitated when hydrochloric ' 
acid or a soluble chloride is added to the solution of a lead salt. 
It is a white solid, only slightly soluble in cold water, but dis- 
solving in thirty-three times its weight of boiling water : when 
the hot solution cools, the chloride separates in brilliant anhydrous 
needles. It is employed in the manufacture of several yellow 
colors, which are oxychlorides of lead, or mixtures of the chloride 
and oxide. 

456. Lead Iodide, Pbl 2 , is deposited as a yellow precipitate 
when potassium iodide is added to the solution of a lead salt. It 



LEAD OXIDES. 275 

is almost insoluble in cold water, but dissolves in a little less than 
two hundred times its weight of boiling water, from which it 
separates on cooling in beautiful golden-yellow scales. 

457. Lead Monoxide, PbO. — This body is produced by the 
direct oxidation of melted lead by the air. It is an accessory prod- 
uct in the cupellation of lead for the extraction of silver (§ 426). 
It is known in commerce by the names massicot and litharge : 
massicot is a yellow, amorphous powder ; by fusing this powder 
and pulverizing the resulting mass, litharge is obtained as reddish- 
yellow, crystalline scales. Lead monoxide is slightly soluble in 
water, and will restore the blue color to reddened litmus. It melts 
at a red heat, but cannot be melted in vessels of glass, porcelain, 
or clay, because it combines with silica and forms a very fusible 
silicate, so destroying the vessel. It is readily reduced by char- 
coal and by hydrogen. It is used in the manufacture of the salts 
of lead : when it is boiled with linseed oil, the latter acquires the 
property of quickly drying or hardening when exposed to the air. 

When the solution of a lead salt is treated with an alkaline 
hydroxide, lead hydroxide, Pb(OH) 2 , is thrown down as a white 
precipitate, soluble in an excess of the alkaline hydroxide. 

458. Lead Dioxide, PbO 2 , is obtained by treating red lead 
with nitric acid. Red lead is a combination of the monoxide and 
dioxide, and the nitric- acid dissolves out the monoxide, forming 
lead nitrate, which is soluble and can be washed out, while the 
dioxide remains as a brown powder. It is not soluble in water, 
and by the action of heat is decomposed into lead monoxide and 
oxygen. It is a very energetic oxidizing agent : a little sulphur 
may be ignited by rubbing it in a mortar with some lead dioxide. 
It absorbs sulphur dioxide, forming lead sulphate ; with hydro- 
chloric acid it forms lead chloride and water, while chlorine is dis- 
engaged. 

PbO 2 + 4HC1 = PbCP + 2H20 + CI 2 

459. Red Lead, (PbO) 2 Pb0 2 .— This body is prepared by 
heating massicot to 300° in furnaces so arranged that it is freely 
exposed to a current of air ; oxygen is then absorbed, and a beau- 
tiful red powder, called minium, or red lead, is formed. It is 



276 LESSONS IN CHEMISTRY. 

plumboso-plumbic oxide, but the proportions of the monoxide 
and dioxide which it contains are not constant, though usually 
responding to the formula given. When heated,' its color darkens ; 
at a red heat it loses part of its oxygen and is converted into 
the monoxide. Red lead is employed as a pigment, and in the 
manufacture of flint-glass, of which the brilliancy and refractive 
power are due to silicate of lead. Mixed into a paste with linseed 
oil, it forms an excellent cement. 

460. Lead Sulphide, PbS. — This compound is the mineral 
galena, which is found in cubical crystals of a bluish-gray color 
and metallic appearance. Its density is 7.58 ; it is much harder 
than lead, and rather brittle. It melts when heated to redness, 
and in contact with air is then oxidized to oxide and sulphate. 
It is converted into lead chloride by boiling with hydrochloric 
acid, hydrogen sulphide being disengaged. Boiling nitric acid 
converts it into lead sulphate. 

461. Tests for Lead. — With the exception of the nitrate 
and acetate, none of the more common lead salts are very soluble. 
Those which are soluble have a sweet and somewhat astringent 
taste. Hydrogen sulphide forms in them a black precipitate of 
lead sulphide : potassium and sodium hydroxides and ammonia pro- 
duce white precipitates, which are soluble in an excess of either 
of the first two reagents. Sulphuric acid yields a white precipi- 
tate even in the most dilute solutions. Hydrochloric acid throws 
down white lead chloride, unless the solution be too dilute ; this 
precipitate is dissolved by boiling, and, on cooling again, sepa- 
rates in crystals. Potassium chromate precipitates yellow lead 
chromate, which is soluble in the alkaline hydroxides. 

If a salt of lead be mixed with sodium carbonate, and heated 
on a piece of charcoal in the inner flame of a blow-pipe, a small 
bead of metallic lead is obtained, and the softness of the bead 
indicates the nature of the metal. 




MAGNESIUM. 277 

LESSON LIII. 

MAGNESIUM.— ZINC— CADMIUM. 

462. These three metals form a natural group, to which belongs also a 
fourth, glucinium, of which the silicate constitutes part of the mineral beryl 
and the green precious stone emerald. They are diatomic metals. 

463. Magnesium, Mg = 24. — This element occurs in nature 
as carbonate in the mineral magaesite, as sulphate in kieserite, and 
as silicate in serpentine and soapstone. The metal is obtained by 
heating its chloride with sodium in an iron crucible, a mixture of 
common salt and calcium fluoride being added as a flux. The so- 
dium is converted into sodium chloride, and the magnesium sepa- 
rates in little globules diffused through the molten mixture, which 
is constantly stirred. When perfectly cold, the mass is broken up, 
and the globules of magnesium are removed and the metal puri- 
fied by distillation in a current of hydrogen. It is now manu- 
factured by decomposing fused carnallite, MgCP.KCl, by an 
electric current. 

Magnesium is a bluish-white metal ; its surface, which is not 
very brilliant, soon tarnishes in the air. Its density is about 
1.75. It is both ductile and malleable, and is ordinarily rolled 
into ribbon or drawn into wire. It does not decompose water at 
ordinary temperatures, but it acts slightly at the temperature of 
boiling. It melts at 500°, and if exposed to the air takes fire 
and burns with great brilliancy. The light of burning mag- 
nesium is very bright, and lamps are constructed in which the 
ribbon is gradually supplied by clock-work. Such lamps are em- 
ployed in photographing the interior of caves and other dark 
localities. The product of the combustion is magnesium oxide 
MgO. Magnesium combines directly with nitrogen. 

464. Magnesium Chloride, MgCl 2 . — When magnesium or 
its oxide or carbonate is dissolved in hydrochloric acid, and the 
solution is concentrated, crystals of magnesium chloride, with six 
molecules of water of crystallization, are obtained. These crystals 



278 LESSONS IN CHEMISTRY. 

cannot be rendered anhydrous, and their solution cannot be evap- 
orated to dryness, for they decompose into hydrochloric acid and 
magnesium oxide. 

MgCl* + H20 = MgO + 2HC1 

Anhydrous magnesium chloride is prepared by dissolving the 
oxide or carbonate in hydrochloric acid, and adding two molecules 
of ammonium chloride for every atom of magnesium. This so- 
lution may be evaporated to dryness, and leaves an anhydrous 
double chloride of magnesium and ammonium. The double salt 
is heated in a clay crucible until all of the ammonium chloride is 
driven off, while the magnesium chloride remains in a state of 
fusion ; on cooling, it solidifies to a pearly-white mass. In this 
form it is used for the manufacture of magnesium. It is very 
soluble in water, but from the solution only the hydrated crystals 
can be obtained. 

465. Magnesium Oxide, MgO. — This is the calcined mag- 
nesia of the pharmacies. It is made by calcining magnesium 
carbonate, or the mixture of hydrate and carbonate commonly 
called magnesia alba. It is a tasteless white powder, infusible 
except in the electric furnace. It is insoluble in water, but com- 
bines with that liquid, forming magnesium hydroxide, Mg(OH) 2 , a 
substance which restores the blue color to reddened litmus. This 
same hydrate is precipitated when an alkaline hydroxide is added 
to the solution of a magnesium salt. 

466. Tests for Magnesium. — Neither hydrogen sulphide nor 
ammonium sulphide occasions any precipitate in magnesium solu- 
tions. Sodium carbonate throws down a white, flocculent precipi- 
tate of the hydrated carbonate, which when dried in the air con- 
stitutes white magnesia. Potassium and sodium hydroxides yield 
white precipitates of the hydroxide, as does also ammonia unless 
the solution be acid or contain ammonium chloride. Sodium phos- 
phate with a few drops of ammonia produces a white, crystalline 
precipitate of magnesium ammonium phosphate, Mg(NH 4 )P0 4 . 

467. Zinc, Zn = 65. — The ores from which zinc is obtained are 
the carbonate, which is called smithsonite, and the sulphide, called 
blende. These minerals are broken up and roasted in furnaces 



ZINC. 



27<J 



much resembling limekilns. At the temperature of the roast- 
ing, which is a dull red heat, the carbonate loses carbon dioxide 
and the water which it usually contains, and is converted into zinc 
oxide : the sulphide is also oxidized by roasting, sulphur dioxide 
being disengaged. The zinc oxide so obtained is mixed with char- 
coal and heated for about twenty-four hours to a high temperature 
in clay or iron vessels : carbon monoxide is disengaged, while the 
zinc volatilizes and is condensed in suitable apparatus. Various 
processes of distillation are employed : we need only consider the 
two which are generally used. In the Belgian process, the mix- 
ture of zinc oxide and charcoal is introduced into clay tubes, closed 
at one end, and inserted in an inclined position in the walls of the 
furnace ; to the open and exterior end of each tube is adapted a 
bulged pipe, in which the zinc vapor condenses and the metal col- 
lects. In order that no air may enter the tubes and oxidize the 
zinc, a sheet-iron noz- 
zle, having a hole for 
the exit of the gases, 
is passed over the ex- 
tremity of this con- 
denser (Fig. 114). 
The tubes are usu- 
ally about a metre in 
length, and twenty 
centimetres in inte- 
rior diameter. A large number of them are placed in parallel 
rows in the same furnace: when all the zinc has distilled, the 
receivers containing it are removed, and a fresh charge of roasted 
ore and charcoal is introduced into the tubes. 

In the Silesian process, the retorts are arched, and very similar 
in form to those employed in the manufacture of illuminating gas 
from coal. 

468. At present, the furnace used in the reduction of zinc by both the Bel- 
gian and Silesian methods is that known as the Siemens regenerative furnace, 
which effects a great saving of fuel. In this arrangement, the coal is fed grad- 
ually to the grate of a peculiar fire-box, called the generator, and the admis- 
sion of air is there so regulated that as much carbon monoxide as possible may 




Fig. 114. 



280 



LESSONS IN CHEMISTRY. 



be produced by an imperfect combustion ; in addition, the ashes below the 
grate are kept moist, and the steam passing into the fire reacts with the hot 
carbon, producing hydrogen and carbon monoxide ($ 232) ; the highly-heated 
gas is led through a chamber filled with fire-bricks, which become very hot ; 
by a system of dampers, the gases are then directed through another similar 
chamber, while air is admitted to that which has been heated ; the heated air 
from the one, and the heated gas from the other, are then brought in contact 
where it is desired that the greatest temperature shall be produced by the per- 
fect combustion of the gases. The heat of the waste products of combustion 
is applied to heating other fire-brick chambers, which will afterwards serve for 

the admission of 
air, as these regen- 
erators, as they are 
called, are cooled 
by the entering 
air. Figure 115 
represents the fire- 
brick chambers of 
a Siemens furnace 
applied to the Si- 
lesian zinc process. 
The two chambers 
on each side serve 
alternately, one 
for the entrance 
of air, and one for 
the gas from the 
generator, while 
the other two serve 
for the exit of the 
products of com- 
bustion. The 
heated air and gas 
from A and A' 
come in contact in 
the space B, and the flames play through openings in the floor above which 
are the clay retorts. The heated products of combustion pass over the retorts 
in another similar chamber, C, and from above downwards through other fire- 
brick chambers, D and D'. The dampers allow the direction of the current of 
gas and air to be reversed from A A' to D D' as often as necessary, and in 
practice it is so changed about once every hour. 

Zinc must usually be purified before it is sent into commerce, 
and the most harmful impurity is lead, for it impairs the mallea- 
bility of the zinc. The lead is separated in great part by melting 




Fig. 115. 



zinc. 281 

the zinc in moulds which are slightly inclined and have a cavity at 
the lower end : in this the greater part of the lead collects by 
reason of its greater density, and may be broken from the cooled 
ingot. Commercial zinc usually contains small quantities of iron, 
copper, lead, cadmium, and sometimes arsenic. Sheet zinc is the 
purest. 

Zinc is a bluish-white metal, capable of taking a high lustre. 
Its density varies from 6.86, that of the cast metal, to 7, that of 
the rolled. Pure zinc may be hammered into sheets, or drawn 
into wire at ordinary temperatures, but commercial zinc must be 
rolled at about 150°. It again becomes brittle at 200°, and may 
readily be pulverized in a mortar heated to that temperature. It 
melts at 410°, and distils at about 1000°. It is unaltered by dry 
air, but in moist air its surface becomes dull from the formation 
of a film of hydrated carbonate, which protects the metal from 
further action. 

When it is heated to redness in the air, it takes fire and burns 
with a bluish flame, giving off clouds of white zinc oxide, ZnO. 
Fine zinc shavings may be lighted by a match, and burn brilliantly 
in the air. If some zinc be heated to redness in a ladle or cruci- 
ble, and pieces of potassium nitrate be thrown in, the oxygen of 
the decomposing nitre energetically oxidizes the metal. 

Zinc is dissolved by hydrochloric, sulphuric, and nitric acids, 
and by boiling solutions of potassium and sodium hydroxides. In 
the latter case, hydrogen is disengaged and an alkaline zincate is 
formed, a compound in which zinc oxide appears to act as an acid 
radical. We have already studied the action of the acids on zinc. 

Zinc is employed in the manufacture of galvanized iron, which 
is made by dipping carefully cleaned iron objects into melted zinc; 
brass, which is an alloy of copper and zinc ; the plates of voltaic 
batteries ; and for the preparation of zinc white, which is zinc 
oxide. 

469. Zinc Chloride, ZnCl 2 , may be formed by the direct 
union of zinc and chlorine, a union which takes place brilliantly 
when fine zinc shavings are thrown into a jar of chlorine. It is 
prepared by dissolving zinc in hydrochloric acid. It forms deli- 



282 LESSONS IN CHEMISTRY. 

quescent crystals containing one molecule of water of crystalliza- 
tion, which is expelled by heat, and the anhydrous salt fuses at 
250°. The latter is very deliquescent, and is an energetic dehy- 
drating agent. It is employed as a caustic in surgery. Zinc 
chloride is very soluble in water, and its solution, to which a little 
free hydrochloric acid and some ammonium chloride have been 
added, is an excellent soldering liquid, for moistening the surface 
of iron, zinc, copper, and brass articles before soldering. 

470. Zinc Oxide, ZnO, is prepared on a large scale by heating 
zinc in large muffles in which its vapor may come freely in con- 
tact with air. The product is stirred up with water; the heavier 
particles of unaltered zinc sink to the bottom, while the zinc oxide 
remains suspended in the creamy liquid which is rapidly poured 
off and allowed to settle. The separation of fine powders by this 
method is called elutriation. 

Zinc oxide is a white powder, insoluble in water. It is em- 
ployed as a substitute for white lead in painting localities exposed 
to hydrogen sulphide, which would blacken a lead pigment. 

When an alkaline hydroxide is added to the solution of a zinc 
salt, zinc hydroxide, Zn(OH) 2 , is thrown down as a white precipitate. 

ZnSO* + 2KOH = K*SO* + Zn(OH)2 
This precipitate is soluble in an excess of the alkaline hydroxide. 

471.. Zinc Sulphide, ZnS. — This compound is found native 
as zinc blende, a mineral usually having a more or less intense 
brown color, due to the presence of a certain proportion of iron. 
When ammonium sulphide is added to the perfectly neutral solu- 
tion of a zinc salt, a white precipitate of hydrated zinc sulphide is 
formed. 

472. Tests for Zinc. — Neutral solutions of zinc salts are 
precipitated white by hydrogen sulphide ; the precipitate is not 
formed if free mineral acid be present. Ammonium sulphide 
produces a characteristic white precipitate of zinc sulphide. The 
alkaline hydroxides and ammonia-water yield white zinc hydrate, 
soluble in an excess of the reagent. Potassium ferrocyanide 
throws down a white precipitate of zinc ferrocyanide. The salts 
of zinc are poisonous. 



CADMIUM. 283 

473. Cadmium, Cd = 112. — This metal occurs associated with zinc in both 
blende and calamine. It is reduced with the zinc, and, being more volatile 
than the latter, distils during the early part of the operation. During the 
first few hours of the reduction of many zinc ores, a brown powder, called 
cadmies, collects in the receivers attached to the retorts. This dust contains a 
large proportion of cadmium oxide, and when distilled with charcoal powder 
yields an alloy of zinc and cadmium. The latter metal is purified by dis- 
solving the alloy in dilute sulphuric acid, precipitating cadmium sulphide by 
passing hydrogen sulphide through the acid liquid, dissolving the sulphide in 
hydrochloric acid, and adding ammonium carbonate. Cadmium carbonate is 
precipitated; this is collected, dried, and roasted, and the cadmium oxide ob- 
tained is distilled with charcoal powder. 

Cadmium has a white color and a brilliant lustre, which soon becomes dull 
in moist air. Its density is 8.60. It melts at 320° and boils at S60°. Hydro- 
chloric and sulphuric acids dissolve it rapidly, disengaging hydrogen. 

474. Cadmium Iodide, Cdl 2 , is made by digesting cadmium filings and 
iodine in water. On evaporating the solution, beautiful transparent and col- 
orless hexagonal prisms of cadmium iodide are deposited. It is used in pho- 
tography. 

475. Cadmium Oxide, CdO, is obtained as a yellowish-brown powder by 
roasting either cadmium nitrate or cadmium carbonate. It is reduced by hy- 
drogen and carbon at lower temperatures than those required for the corre- 
sponding reductions of zinc oxide. 

476. Cadmium Sulphide, CdS, is found in nature as greenochite in brilliant 
yellow, hexagonal prisms. It is precipitated as an amorphous yellow powder 
by the action of hydrogen sulphide on solutions of cadmium salts. It is em- 
ployed as a pigment by artists. 

477. Tests for Cadmium. — Potassium and sodium hydroxides and ammonia- 
water give white precipitates of cadmium hydroxide ; only that formed by am- 
monia is soluble in an excess of the reagent. Hydrogen sulphide throws 
down a characteristic yellow precipitate of cadmium sulphide, even in acid 
solutions. Potassium ferrocyanide gives a yellowish-white precipitate of 
cadmium ferrocyanide. 



LESSON LIV. 

COPPER Cu = 63.1. 

478. Large deposits of metallic copper exist on the shores of 
Lake Superior, the metal being sometimes found in crystals, some- 
times in irregular and grotesque masses. The more common 



284 



LESSONS IN CHEMISTRY. 



copper ores are cuprous sulphide, called chalcocite, and copper 
pyrites, a compound of cuprous sulphide and iron sulphide. 
This metal is also found as cuprous oxide, cupric oxide, and cupric 
carbonate. 

Pure copper ores — those containing only the oxide, carbonate, 
or sulphide of copper, and very little of other metals — are easily 
reduced : the sulphide is first converted into oxide by roasting, 
and the ores are then heated with charcoal in a somewhat con- 
ical furnace. The reduction of copper pyrites is more difficult, 
especially if, as is often the case, this mineral be mixed with the 
sulphides of antimony, arsenic, zinc, etc. If such ore contains a 
large proportion of copper, it may be worked by a dry process; but 
if only a small percentage of copper is present, a method of solu- 
tion is adopted. In the dry process, the ore is first roasted by 
being fed from hoppers on to the hearth of a reverberatory fur- 
nace (Fig. 116), where it is swept by the flame of a fire. Part of 




the sulphur is so converted into sulphurous oxide, which may be 
used for the manufacture of sulphuric acid, while the iron and 
copper of the pvrites are partially converted into oxide and sul- 
phate. A quantity of sand and silicate of iron from a subsequent 
stage of the operation is then added, and the mass is transferred 
either to rotating cylindrical furnaces or to reverberatory furnaces 
with deep hearths, where it can be strongly heated. The un- 



copper. 285 

altered ferrous sulphide remaining in the roasted mass then 
reacts with the cupric oxide formed, and the result is cuprous 
sulphide and ferrous oxide. The latter unites with the silica, 
forming ferrous silicate, which is very fusible, and is drawn off 
as slag, while cuprous sulphide containing some iron sulphide 
collects on the hearth of the furnace. This product, which is 
called copper matte, is broken up, and repeatedly roasted until 
nearly all the sulphur is expelled, and a considerable propor- 
tion of the copper is reduced to the metallic state ; the more 
oxidizable foreign metals present become oxidized, and, on the 
addition of silicious matters, are converted into fusible silicates 
by an increased temperature. The black copper so obtained 
contains from 90 to 94 per cent, of copper, the remainder being 
lead, iron, sulphur, arsenic, etc. 

The extraction of copper from the matte is now frequently effected by 
forcing air through the molten material in a Bessemer converter (p. 312). 
The iron and part of the copper are oxidized, while sulphur dioxide is dis- 
engaged. The remaining sulphide of copper then reacts with the cuprous 
oxide so as to yield the metal, and the ferrous oxide forms a slag with the 
silicious lining of the converter. 

The crude'metal may be refined by melting it on the hearth 
of a reverberatory furnace and exposing it to an oxidizing 
atmosphere. The impurities are thus completely oxidized and 
removed, either by volatilizing or by forming a slag. After 
removing the latter, some of the copper is allowed to oxidize 
to destroy the last traces of sulphur. The cuprous oxide is 
finally reduced by stirring the molten metal with poles of green 
wood ; the combustible gases formed by the action of the high 
temperature on the wood completely deoxidize the copper, and 
the cold metal is red and soft. 

Another mode of refining, known as the electrolytic process, 
consists in dissolving the impure copper in an acid, and then 
precipitating the metal by means of the electric current. In an 
acid bath of copper sulphate solution, plates of the crude metal, 
constituting the anodes, alternate with thin copper plates which 
serve as cathodes. The current causes the anodes to dissolve, 



286 LESSONS IN CHEMISTRY. 

while it deposits an equal amount of copper upon the cathodes : 
some of the impurities, including the silver and gold, remain 
undissolved, forming the valuable anode mud ; the others are 
retained in the bath. A very pure product is thus obtained. 

In the Lake Superior district, where enormous quantities of 
native copper are found, the metal is separated from the rock 
by mechanical means, and subsequently subjected to a refining 
process analogous to that first described. " Lake Copper" in- 
cludes the most valued brands of the copper of commerce : it 
is remarkable for its high electric conductivity. 

Large quantities of copper are extracted also by the wet 
process, particularly from the burnt pyrites of the sulphuric 
acid works. In furnaces of peculiar construction the ore mixed 
with common salt is roasted, whereby the copper is converted 
into cupric chloride. This is extracted with water, and the 
copper then precipitated in the metallic state by placing scrap- 
iron in the solution. 

CuCl* + Fe = FeCl 2 + Cu 

Cement copper is obtained by precipitating with iron the cop- 
per sulphate solutions occurring in certain localities. 

Copper has a red color and a brilliant lustre. Its density is 
about 8.9. It is exceedingly ductile, malleable, and tenacious, 
and one of the best conductors of heat and electricity. It 
melts at about 1100°, and may be crystallized either by fusion 
or by electrolysis of solutions of its salts. It vaporizes in the 
electric arc, and in the oxyhydrogen flame. 

Copper is unaltered by cold dry air, but by moist air it is 
gradually converted into a hydrocarbonate, which appears in green 
spots on the surface of the metal. This is the substance ordinarily 
called verdigris (see § 333). 

At a temperature about redness, copper combines directly with 
oxygen, forming either cupric oxide, CuO, or cuprous oxide, 
Cu 2 0, according to the access of air. When copper acetate is 
strongly heated in a hard glass tube, it is entirely decomposed, 
and a residue of finely-divided copper is obtained. If this be 
turned out and heated at one point by a lighted match, a black 



copper. 287 

spot appears and rapidly spreads over the entire mass, which is so 
converted into cupric oxide. 

We have already studied the action of sulphuric and nitric 
acids on copper. Hydrochloric acid attacks it only when boiling, 
and then but slowly, evolving hydrogen, and forming cuprous 
chloride, Cu 2 CP. 

Ammonia in presence of oxygen exerts a curious action on copper. We in- 
troduce some copper clippings and a little ammonia into a bottle, which we ' 
tightly cork and then agitate for a few minutes. The liquid becomes blue, 
and if we invert the bottle and open it with its mouth under water, the latter 
will rise in the bottle, showing that part of the air has been absorbed. It is 
the oxygen which is absorbed, and the blue liquid contains copper nitrite and 
ammoniacal cupric oxide, both of which are soluble in ammonia. This liquid 
is capable of dissolving cotton, linen, paper, and other forms of cellulose. 

Copper is used for the manufacture of boilers, stills, con- 
densing apparatus, and other utensils for the laboratory, manu- 
factory, and kitchen. As wire, it is extensively used for electrical 
purposes. In sheets, it serves for sheathing ships, and some- 
times for roofing. It constitutes part of many alloys, among 
which are brass, containing from 65 to 90 per cent, copper, the 
remainder being zinc ; a large proportion of copper gives a red 
color to the brass ; these metals are melted together in crucibles, 
the zinc being added after the copper is fused. Bronze contains 
from 93 to 95 per cent, of copper, the remainder being tin, with 
sometimes 1 per cent, of zinc. Gun metal is about 91 per cent, 
copper and 9 per cent. tin. Bell-metal and the very white specu- 
lum metal contain respectively 78 and 67 per cent, of copper, the 
remainder being tin. German silver is an alloy of copper, zinc, 
and nickel. The United States cents contain 95 per cent, of 
copper, 2.5 per cent, of zinc, and 2.5 per cent, of tin. 

479. Copper forms two series of compounds. It is a diatomic 
element, and in the cuprous compounds two atoms of copper form 
a diatomic couple, Cu-Cu, which replaces two atoms of hydrogen 
in the acids. In the cupric compounds, a single diatomic atom 
of copper replaces two atoms of hydrogen. 

480. Cuprous Chloride, Cu 2 C1 2 , may be made by boiling a 
solution of cupric chloride with copper, or by boiling copper with 
hydrochloric acid and adding a little nitric acid from time to time ; 



288 LESSONS IN CHEMISTRY. 

in the latter case, cupric chloride is formed, and is at once re- 
duced by the metallic copper present. On adding water to the 
brown liquid so obtained, cuprous chloride is thrown down as a 
white crystalline precipitate. It is insoluble in water, but dis- 
solves in ammonia, forming a colorless solution which absorbs 
oxygen and becomes blue on exposure to air. It also dissolves 
in hydrochloric acid, and both of these solutions are capable of 
absorbing a large volume of carbon monoxide. 

481. Cupric Chloride, CuCl 2 , is obtained when cupric oxide 
is boiled in hydrochloric acid. When the green solution is suf- 
ficiently concentrated, it deposits beautiful bluish-green crystals 
of cupric chloride, with two molecules of water of crystallization. 

A hydrated compound of cupric chloride and cupric oxide 
is met with in the beautiful green mineral atacamite. 

482. Cuprous Oxide, Cu 2 0. — This substance constitutes the 
beautiful mineral cuprite which occurs in red octahedra and 
cubes. It may be made by boiling glucose with a solution of 
cupric acetate, and is then thrown down as a bright red crys- 
talline precipitate. If heated in contact with air, it is con- 
verted into cupric oxide. The color of red copper glass is due 
to the presence of this compound. 

483. Cupric Oxide, CuO. — When cupric nitrate is strongly 
heated, it yields a fine black powder of cupric oxide. This com- 
pound is usually prepared by heating metallic copper to redness in 
vessels through which air is blown or drawn. The copper then 
absorbs oxygen, and is converted into hard and compact cupric 
oxide. This substance is reduced by both hydrogen and charcoal 
at temperatures below redness, water or carbon dioxide being 
formed. It communicates a green color to glass, and is used for 
that purpose. In the laboratory it is of great value in the analysis 
of carbon compounds. 

When potassium or sodium hydroxide is added to the solution of 
a cupric salt, cupric hydroxide, Cu(OH) 2 , is formed as a pale-blue 
precipitate. When the liquid containing this hydroxide is boiled, 
the precipitate turns black, for it is converted into cupric oxide 
and water, even when surrounded by liquid. 



COMPOUNDS OF COPPER. 289 

484. Cuprous Sulphide, Cu 2 S. — This is the mineral chalco- 
cite. It may be obtained as a black, brittle, crystalline mass by 
fusing together sulphur and copper, or by burning copper in vapor 
of sulphur. 

485. Cupric Sulphide, CuS, is thrown down as a brownish- 
black precipitate by the action of hydrogen sulphide on cupric 
solutions. When heated, it loses sulphur, and is converted into 
cuprous sulphide. 

486. Carbonates of Copper. — When a solution of cupric 
sulphate is treated with sodium carbonate, carbon dioxide is dis- 
engaged, and a bluish-green precipitate is thrown down ; when 
washed with warm water, its color becomes green ; it is a com- 
pound of cupric hydroxide and cupric carbonate, containing 
CuC0 3 .Cu(OH) s . The beautiful green mineral malachite, which 
when polished displays veins of variegated tints, is a compound 
having the same composition. Azurite, a mineral found in fine 
blue crystals, is a compound of two molecules of cupric carbonate 
with one of cupric hydroxide, Cu(OHy.2CuC0 3 . 

487. Tests for Copper. — The salts of copper have either blue 
or green colors. Both hydrogen sulphide and ammonium sulphide 
throw down brownish-black precipitates. The alkaline hydroxides 
precipitate pale-blue cupric hydroxide, insoluble in an excess of the 
reagent. Ammonia also produces a pale-blue precipitate, but 
this dissolves when an excess of ammonia is added, yielding a 
magnificent blue solution of an ammonio-cupric salt. 

Potassium ferrocyanide produces a mahogany-brown precipitate 
of cupric ferrocyanide, and the test is exceedingly delicate and 
characteristic. A clean piece of iron, as a needle or knife-blade, 
dipped in a cupric solution, quickly becomes covered with a red 
layer of metallic copper : this test is conclusive. 



19 



290 



LESSONS IN CHEMISTRY. 



LESSON LV. 



MERCURY. Hg = 200. 



488. Mercury is found in small quantity in the metallic state, 
but its principal ore is the sulphide, which constitutes the mineral 
cinnabar. It is especially abundant in Spain and on the Pacific 
slope. 

The reduction of cinnabar is a simple operation : it is broken 
up and roasted in a current of air, the sulphur being expelled as 
sulphur dioxide, while mercury distils. Very little improvement 
has been effected in the furnaces during hundreds of years ; the 
mercury vapor is sometimes condensed by being passed through a 
long series of clay pipes, sometimes by being directed through a 
number of chambers containing a layer of water, by which the 
gases are cooled (Fig. 117). The mercury is then filtered through 




Fig. 117. 

closely-woven canvas, and is usually transported in iron bottles, 
each bottle holding about sixty pounds. 

Mercury is liquid at ordinary temperatures : it freezes at — 40°, 
and boils at 357°. Its density at 0° is about 13.6. 



; 



MERCURY. 291 

The density of mercury vapor compared with that of hydrogen is 100 : its 
atomic weight is 200, as is shown by the vapor-densities of its volatile com- 
pounds. Then if equal volumes of gases contain equal numbers of molecules, 
and if the molecule of hydrogen contain two atoms, the molecule of mercury 
vapor must consist of a single atom. This is the case also with zinc, cadmium, 
argon, helium, and a few other elements. 

Mercury is unaffected by the air at ordinary temperatures, but 
at 300° it absorbs oxygen and is converted into red mercuric 
oxide. It combines directly, and in the cold, with chlorine, bro- 
mine, and iodine, and with sulphur by the aid of a gentle heat, or 
if the sulphur be finely divided. Mercury is not dissolved by 
hydrochloric acid : boiling sulphuric acid converts it into mercuric 
sulphate, sulphur dioxide being disengaged. Nitric acid dissolves 
it, emitting red vapors, and forming mercurous nitrate if the re- 
action take place in the cold, or mercuric nitrate if the acid be 
boiling. 

Mercury is used for filling thermometers, barometers, and press- 
ure-gauges ; for silvering ordinary mirrors, which are coated with 
tin foil amalgamated with mercury ; for the extraction of silver 
and gold from their ores ; and for the preparation of various 
amalgams. 

The mercury of commerce is rarely pure ; it contains small 
quantities of lead, copper, tin, and sometimes bismuth. Its ap- 
proximate purity may be determined by allowing a few drops to 
fall on a clean piece of paper or porcelain ; pure mercury will 
then break up into small globules which are perfectly round, and 
move about freely when the surface on which they rest is inclined, 
but mercury containing other metals forms globules that are drawn 
out to a tail, and that do not move so readily. The surface of 
pure mercury is perfectly brilliant, but when impure the metal 
has a tarnished appearance. It may be purified by treating the 
metal in a finely divided state with very dilute nitric acid, or 
by distillation in vacuo. 

Like copper, mercury is diatomic, and forms two series of com- 
pounds, — mercurous compounds, in which two atoms form a di- 
atomic couple, and mercuric compounds, in which two atoms of 
hydrogen are replaced by a single diatomic mercury atom. 



292 LESSONS IN CHEMISTRY. 

489. Mercurous Chloride, Hg 2 Cl 2 . — This compound is the 

well-known medicine calomel. It is made by subliming a mixture 

of mercurous sulphate and common salt. 

Hg2SO* + 2NaCl = Na2SO± + Hg2Cl 2 

Mercurous sulphate. Mercurous chloride. 

The calomel then condenses in appropriate receivers, in dense 
crystalline masses. It is usually resublimed, and its vapors passed 
into jars or chambers filled with steam, where it condenses in an 
impalpable powder. Calomel is precipitated when hydrochloric 
acid is added to the solution of a mercurous salt. 

In masses, calomel occurs in dense, fibrous, crystalline, and 
translucent fragments, colorless when recently prepared, but be- 
coming gray or yellowish by the action of light which partially 
decomposes this compound into mercuric chloride and mercury. 
Its density is about 7.2. It is insoluble in water ; when calomel 
is agitated with water and the liquid filtered, no turbidity should 
be produced in the filtrate by the addition of sodium carbonate 
solution. If mercurous chloride be heated with a solution of 
sodium chloride, it is converted into mercuric chloride, while 
metallic mercury is deposited as a gray powder. 

490. Mercuric Chloride, HgCl 2 , is the poisonous compound 

corrosive sublimate. It is prepared by subliming a mixture of 

common salt and mercuric sulphate, sodium sulphate being formed 

at the same time. 

HgSO* + 2NaCl = N a 2SO± + HgCl 2 
Mercuric sulphate. Mercuric chloride. 

It is also formed by the direct combination of chlorine and 
mercury. 

It forms dense, white or colorless, crystalline masses, having a 
density of 6.5. It melts at 265°, and boils at about 295°. It is 
soluble in nineteen times its weight of cold water, and in much 
less boiling water, from which it separates in anhydrous crystals 
on cooling. It is exceedingly poisonous, and its antidote is white 
of egg, for it forms an insoluble compound with albumen. 

491. Mercurous Iodide, Hg 2 I 2 , is obtained as a green powder 
by rubbing together in a mortar 100 parts of mercury and 63.5 



COMPOUNDS OF MERCURY. 293 

parts of iodine with a few drops of alcohol. By the action of 
light or heat, it is decomposed into mercuric iodide and metallic 
mercury. 

492. Mercuric Iodide, Hgl 2 . — This beautiful compound is 
prepared by mixing potassium iodide with four-fifths its weight 
of mercuric chloride, both in aqueous solution, and thoroughly 
washing the precipitate. 

HgCl 2 + 2KI = Hgl 2 + 2KC1 

If either substance be employed in excess, the precipitate will be 
redissolved. 

So obtained, mercuric iodide forms a dark-red powder, which is 
almost insoluble in water, but dissolves slightly in boiling alcohol, 
and on cooling separates in red, octahedral crystals. 

Mercuric iodide presents a curious case of dimorphism. If a 
little of the red powder be cautiously heated on a sheet of white 
paper on which it is spread out, the red color changes to yellow ; 
the yellow particles are rhombic prisms, and if they be rubbed 
with a glass rod or any hard body, they will reassume the red color 
and their first crystalline form, the octahedron. Mercuric iodide 
melts to a dark-yellow liquid, and volatilizes, condensing in the 
yellow crystals. 

With potassium iodide, mercuric iodide forms a soluble com- 
pound, which may be obtained by dissolving the mercuric iodide 
in solution of potassium iodide. The colorless liquid is called 
Nesslers reagent, and is used in the laboratory as a test for am- 
monia, and compound ammonias, with which it forms a brownish 
cloud or a dense precipitate, according to the proportion of am- 
monia present. 

493. Mercurous Oxide, Hg 2 0. — This substance is obtained 
as a black powder by digesting mercurous chloride in a solution 
of potassium hydroxide. A temperature of 100°, or the prolonged 
action of light, decomposes it into mercuric oxide and mercury. 

494. Mercuric Oxide, HgO, has long been known under the 
name red precipitate. It may be made either by decomposing 
mercuric nitrate by heat until the whole is converted into a red 
powder and no more red vapors are disengaged, or by adding po- 



294 LESSONS IN CHEMISTRY. 

tassium hydroxide to a solution of mercuric chloride and thoroughly 
washing the precipitate. Prepared in the first manner, it forms 
a red, crystalline powder ; obtained by precipitation, it is yellow 
and amorphous, but becomes red when heated. 

Mercuric oxide is insoluble in water : when it is heated, its color 
darkens, and at a temperature of about 400° it is decomposed into 
metallic mercury and oxygen. It is an energetic oxidizing agent. 
In presence of water, it converts chlorine into hypochlorous acid, 
and when dry and quite cold, into hypochlorous oxide. If a mix- 
ture of a little mercuric oxide and sulphur be heated in a test-tube, 
it explodes. 

495. Mercuric Sulphide, HgS. — This is the mineral cin- 
nabar, which is found in hard dense masses, and in transparent 
red crystals. It is manufactured by grinding together the required 
proportions of mercury and sulphur, and subliming the resulting 
black mass. It then forms a dark-red, crystalline solid, having a 
density of 8.12. When strongly heated out of contact with air, 
it volatilizes without melting. When heated in the air, it takes 
fire, and burns with a blue flame, mercury vapor and sulphur 
dioxide being disengaged. 

The fine scarlet pigment vermilion is very finely divided mer- 
curic sulphide, made by grinding for a long time in a mortar a 
mixture of 300 parts of mercury and 114 parts of flowers of 
sulphur : 75 parts of potassium hydroxide dissolved in 400 parts 
of water, are then added, and the grinding is continued, the mortar 
being kept at a temperature of about 45°. When the powder has 
assumed the desired shade, it is quickly washed with hot water, 
and dried. 

496. Tests for Mercury. — Very few of the mercurous salts 
are soluble : in their solutions, hydrochloric acid produces a white 
precipitate of mercurous chloride ; hydrogen sulphide and potassium 
and sodium hydroxides and ammonia produce black precipitates. 

With mercuric salts, hydrogen sulphide and ammonium sulphide 
give black precipitates ; potassium hydrate throws down yellow 
mercuric oxide. If a piece of bright copper be dipped into the 
slightly acid solution of either a mercurous or a mercuric salt, 



: 



BISMUTH. 295 

metallic mercury is quickly deposited on the copper, whose surface 
becomes white and brilliant after a little friction. 

When heated with lime or sodium carbonate in a small glass 
tube, all compounds of mercury yield a sublimate of metallic mer- 
cury, which condenses in the cooler part of the tube in microscopic 
globules. On throwing a fragment of iodine into the still warm 
tube, the globules are changed into yellow or red mercuric iodide. 



LESSON LVI. 

BISMUTH AND GOLD. 

These two metals are triatomic : they form chlorides whose molecules contain 
one atom of metal and three atoms of chlorine. They form trioxides, containing 
two atoms of metal and three of oxygen. Of bismuth we know also a dioxide, 
Bi 2 2 , a tetroxide, Bi 2 4 , and a pentoxide, Bi 2 5 , and gold forms an oxide, 
Au 2 0. 

497. Bismuth, Bi = 206.5. — Bismuth is found in the metallic 
state disseminated in quartz. It is separated from the earthy 
materials, which are called the gangue, by heating the mineral in 
iron tubes which are closed at one end, and arranged in an in- 
clined position in a furnace beyond which the lower and open end 
projects. The bismuth then melts and runs out of the tubes. 
The bismuth thus obtained is never pure, but contains small 
quantities of other metals, and sometimes traces of arsenic and 
sulphur. In order to purify it, it is pulverized and mixed with a 
little potassium nitrate : the mixture is heated to redness in clay 
crucibles ; the impurities, which are more easily oxidized than 
the bismuth, are thus oxidized, and any arsenic present is con- 
verted into potassium arsenate. 

Bismuth is a crystalline, brittle, reddish-white metal. Its 
density is 9.8. It melts at 264°. By allowing a crucible full of 
the molten metal to cool until a crust forms on the surface, and 
then pouring out the liquid interior through a hole made in the 
crust, fine crystals of bismuth may be obtained. These crystals 
become superficially oxidized, and the thin film of oxide imparts 



296 LESSONS IN CHEMISTRY. 

to them all the colors of the rainbow. Bismuth is unaffected by 
cold air, but at a red heat it is burned to bismuth oxide. It dis- 
solves in nitric acid, forming bismuth nitrate, while red vapors are 
disengaged. 

In addition to its use for the preparation of the bismuth com- 
pounds, this metal is employed chiefly for the manufacture of 
certain alloys. Britannia metal contains about one per cent, of 
bismuth. The fusing points of the bismuth alloys are much 
lower than that of bismuth. A mixture known as Wood's alloy 
or fusible metal consists of one or two" parts of cadmium, two 
parts of tin, four of lead, and seven or eight of bismuth. It 
melts between 66° and 71°, according to its composition. An- 
other alloy, known as Arcet's fusible metal, is made by melting 
together eight parts of bismuth, five of lead, and three of tin. 
It melts at 94.5°. 

Bismuth much resembles antimony in many of its chemical re- 
lations ; but we class it among the metals, because it is capable of 
replacing the hydrogen of oxygen acids, so forming well-defined 
salts. 

498. Bismuth Chloride, BiCl 3 . — This compound results from the direct 
union of chlorine and bismuth. When powdered bismuth is sprinkled into 
chlorine, it burns brilliantly, forming the chloride. This substance is pre- 
pared by passing dry chlorine over melted bismuth in a retort so arranged 
that the chloride may collect in a receiver as it distils. It then forms a crys- 
talline deliquescent mass, which is quite soft at ordinary temperatures, being 
very fusible. It is soluble in hydrochloric water, but is decomposed by water, 
hydrochloric acid being formed, while a white powder of bismuth oxychloride, 
BiOCl, is thrown down. 

2BiCl 3 + 2H20 = 4HC1 + 2B10C1 

Bismuth oxychloride constitutes the cosmetic known as pearl-white. 

499. Bismuth Oxide, Bi 2 3 , is obtained as a yellow powder when bismuth 
nitrate is strongly heated. It melts at a red heat, and on cooling solidifies to 
a glassy, yellow mass. It forms a very fusible silicate, and therefore cannot 
be melted in clay crucibles. Bismuth hydroxide, probably Bi(OH) 3 , is thrown 
down as a white powder when bismuth subnitrate is treated with potassium 
hydroxide or ammonia-water. 

500. Bismuth Nitrate, Bi(N0 3 ) 3 . — When bismuth is boiled 
with nitric acid, and the solution is concentrated, large, colorless, 



GOLD. 297 

deliquescent crystals of bismuth nitrate with three molecules of 
water of crystallization are deposited. Since bismuth is triatomic, 
one atom of bismuth will replace the hydrogen in three molecules 
of nitric acid, and combine with the three groups NO 3 . The 
crystals of bismuth nitrate are very soluble in water containing 
free nitric acid; but if the solution be diluted with a large volume 
of water, a pulverulent white precipitate is thrown down. This 
contains (BiO)NO 3 , or BiNO*, and is employed in medicine under 
the name subnitrate of bismuth. A larger quantity may be ob- 
tained by adding very dilute ammonia to the liquid. 

501. Tests for Bismuth. — When solutions of the bismuth 
salts are largely diluted with water, white precipitates of sub-salts 
are thrown down. Hydrogen sulphide and ammonium sulphide 
occasion brown precipitates of bismuth sulphide. The alkaline 
carbonates and hydroxides yield white precipitates, insoluble in an 
excess of the reagent. 

When a bismuth salt is heated with sodium carbonate in the 
inner flame of a blow-pipe, a brittle bead of metallic bismuth is 
obtained. 

502. Gold, Au = 195.7. — Gold is found in the metallic state, 
sometimes in masses called nuggets, but more usually in small 
particles disseminated through quartz rock, or the sand produced 
by the disintegration of the rock. It is sometimes associated with 
silver, copper, lead, and tellurium. The gold is extracted from 
gold-bearing sand by washing the latter in a stream of running 
water in troughs called cradles. By reason of its great density, 
the gold then sinks to the bottom, while the lighter sand is carried 
on with the water. The gold may then sometimes be removed at 
once ; sometimes it is in such small particles that it must be amal- 
gamated with mercury, as will presently be described. Quartz 
rock containing gold is crushed by powerful machinery, and the 
greater part of the earthy matter is removed by washing in vessels 
containing mercury, which forms an amalgam with the gold. Fig. 
118 represents an apparatus which is sometimes employed for 
grinding together the mercury and crushed rock. It consists of 
inclined iron basins, each containing two cast-iron balls : the rock 



298 LESSONS IN CHEMISTRY 

and mercury being introduced into these vessels, a motion of rota- 
tion is communicated by machinery, and by the friction of the 




Fig. 118. 

balls the rock is reduced to an impalpable powder, which is carried 
off by a current of water flowing through the basins, while the 
gold amalgamates with the mercury. From the resulting amal- 
gam the mercury is removed by distillation. 

Very large quantities of gold are now obtained by what is 
known as the cyanide process. It depends upon the solubility 
of metallic gold in a solution of potassium cyanide (see p. 168). 

Au + 2KCN + H 2 = KAu(CN) 2 + KOH + H 

When the crushed rock containing gold in a finely divided 
state is treated with dilute potassium cyanide solution, the 
greater part of the metal is dissolved. It may be precipitated 
by the electric current or by means of metallic zinc. 

2KAu(CN)2 + Zn = K2Zn(CN)± + 2Au 

As extracted from its native rocks or sand, gold is rarely pure. 
It usually contains more or less silver ; this may be removed by 
boiling the metal in nitric acid, which does not affect the gold, 
while it converts the silver into silver nitrate. However, if only 
a small proportion of silver be present, that metal is protected 
by the gold, and it is necessary to melt the alloy with a larger 
proportion of silver before boiling it with nitric acid. The gold 
then remains as a spongy mass. Pure gold may also be obtained 
by adding ferrous sulphate or oxalic acid to a solution of gold 



AURIC CHLORIDE. 299 

chloride ; in this case the gold is thrown down as a dark-brown, 
dull powder, capable of assuming its natural high lustre by 
burnishing. 

The color of gold varies from greenish yellow to a red almost 
as decided as that of copper. Light which has been successively 
reflected from ten surfaces of gold is scarlet. Gold is quite soft, 
and the most malleable and ductile of the metals. Its density is 
19.3 ; it melts at about 1200°, and at a higher temperature emits 
a green vapor. A thin gold-leaf, carefully spread out between 
two plates of glass, allows the passage of a faint green light. 

Gold is not oxidized by air, either moist or dry, or at any tem- 
perature. It is not affected by boiling with nitric, sulphuric, or 
hydrochloric acids. Nitro-hydrochloric acid dissolves it, disen- 
gaging red vapors, and forming a yellow solution of gold tri- 
chloride ; the nitro-hydrochloric acid employed for dissolving gold 
is a mixture of nitric acid with four times its weight of hydro- 
chloric acid. Gold is also attacked by selenic acid, H 2 SeO", by a 
hot mixture of iodic and sulphuric acids, and by a boiling mixture 
of concentrated nitric and sulphuric acids ; from the latter solu- 
tion the gold is again deposited in the metallic state by the addi- 
tion of water. Gold dissolves readily in chlorine-water, in bromine, 
and combines directly with iodine under the influence of light. 

Gold forms two series of compounds, — aurous compounds, in 
which the metal appears to be monatomic, and auric compounds, 
in which it is triatomic. 

503. Auric Chloride. AuCP. — When the solution of gold in 
nitro-hydrochloric acid is evaporated, auric chloride is deposited as 
a dark-red crystalline mass, which is very deliquescent. It is very 
soluble in water and in ether. Its strong solutions are orange 
brown, but the dilute solution is pure yellow. It produces a violet 
stain on the skin ; it is decomposed by the action of heat, and 
more slowly by light, and is reduced by many substances, among 
which are phosphorus, phosphorous and sulphurous acids, oxalic 
acid, and ferrous sulphate. A stick of phosphorus immersed in 
an ethereal solution of auric chloride becomes quickly coated 
with a film of gold. The metal is deposited as a brown powder 



300 LESSONS IN CHEMISTRY. 

when either ferrous sulphate or oxalic acid is added to a solution 
of auric chloride. 

When a solution containing a mixture of stannous and stannic 
chlorides is added to auric chloride, a flocculent, purple precipi- 
tate of uncertain composition, but containing gold, tin, oxygen, and 
hydrogen, is thrown down. This precipitate is known as purple 
of Cassius, and is employed in painting on glass and porcelain. 

When auric chloride is heated to 230°, chlorine is disengaged, 
and an insoluble yellow powder of aurous chloride, AuCl, remains. 

There are two oxides of gold, — aurous oxide, Au 2 0, and auric 
oxide, Au 2 3 . The first is basic, the second forms aurates with 
the metals. When caustic alkalies are fused with gold in contact 
with air, alkaline aurates are formed. 

504. Assaying of Gold. — Gold coin, jewelry, etc., are generally alloyed 
with silver, and sometimes with copper. A weighed quantity of the metal to 
be assayed is first melted with about three times its weight of silver, and the 
resulting button is cupelled in a bone-ash cupel (§ 433). The copper and any 
other base metals present are so converted into oxides, which are absorbed by 
the cupel, and a button containing only gold and silver is obtained. This is 
hammered out into a thin sheet, which is twisted up and boiled in nitric acid ; 
the silver is dissolved, while the gold remains as a spongy mass, which is 
washed, heated to redness, and then weighed. 

The gold coin of the United States contains 90 per cent, of gold, the re- 
mainder being copper. 

505. Gilding. — Silver and copper objects may be gilded by rubbing over 
them an amalgam of gold with eight times its weight of mercury. They are 
then heated under a chimney so arranged that the poisonous mercury vapor 
may be entirely carried off. The dull gilded surface is then rendered brilliant 
by burnishing. A thin film of gold is deposited on copper objects when they 
are dipped into a hot solution of auric chloride with sodium carbonate and 
sodium phosphate. 

Gilding is best accomplished by connecting the objects to be gilded with the 
zinc pole of a voltaic battery, and immersing them in a solution obtained by 
boiling auric chloride with potassium cyanide. The positive pole of the bat- 
tery is connected with a plate of gold immersed in the same liquid. 

The rare elements, gallium, indium, and thallium, which were discovered by 
the aid of the spectroscope, are related to gold and bismuth in the general con- 
stitution of their compounds. Traces of indium and of gallium exist in many 
zinc blendes, while small quantities of thallium occur in certain iron pyrites, 
and the metal is obtained from the dust which collects in the flues of sulphuric 
acid works when these pyrites are burned for the production of sulphur dioxide. 



ALUMINIUM. 301 

LESSON LVII. 

ALUMINIUM. Al = 27. 

506. This is one of the most abundant elements, but it is 
found only in combination. It may be obtained in various 
ways, but is now generally manufactured by decomposing the 
oxide, APO 3 , dissolved in a bath of cryolite, by powerful electrie 
currents. The cryolite is melted in an electric furnace in which 
the carbon lining forms the cathode and carbon rods the anode. 
Pure aluminium oxide, or corundum, is now added : it dissolves 
and is reduced, aluminium collecting on the hearth, while the 
oxygen forms carbonic oxide with the positive carbon. This 
process is continuous. 

Aluminium was formerly obtained by heating a mixture of 

the double chloride of aluminium and sodium with metallic 

sodium. 

AlC-l 8 ,XaCl + 3Xa = 4XaCl + Al 

Aluminium is a tin-white metal, capable of being highly 
polished. It is very ductile and malleable, and also very sonorous. 
It is a good conductor of heat and electricity. Its density is 2.56 ; 
it is therefore as light as glass and porcelain. It melts at about 
750°. It is unaltered by the air at ordinary temperatures, but 
when melted absorbs oxj-gen and is converted into aluminium 
oxide. It is hardly affected by either nitric or sulphuric acid, 
but dissolves readily in hydrochloric acid, disengaging hydrogen 
and forming aluminium chloride. It is also dissolved by boiling 
solutions of the alkaline hydrates, hydrogen being set free, while 
alkaline aluminates are formed. 

The great tenacity of aluminium, and its lightness and un- 
changeableness in the air, render it an exceedingly valuable 
metal. It is employed for scientific instruments, cooking uten- 
sils, ornaments, and as a reducing agent in certain metallurgical 
operations. 



302 



LESSONS IN CHEMISTRY. 



Aluminium bronzes are alloys of copper and aluminium. One 
of them, containing ten per cent, of aluminium, possesses ex- 
traordinary tenacity, and does not tarnish in the air. Its color 
resembles that of gold. 

507. Aluminium Chloride, A1C1 3 . — When aluminium or its 
hydroxide is dissolved in hydrochloric acid, a solution of alu- 
minium chloride is obtained, but this solution cannot be evapo- 
rated to dryness without decomposing into aluminium oxide 
and hydrochloric acid. 

2A1C1 3 + 3H20 = A1 2 3 + 6HC1 

Solid aluminium chloride is formed by passing chlorine gas over 
a red-hot mixture of aluminium oxide and charcoal, which has 
been made into small balls with a little oil, and then calcined in a 
crucible. These balls are put in a clay tube or retort, which is 
heated to bright redness, and dry chlorine is then passed through 
(Fig. 119). Carbon monoxide and aluminium chloride are 




Fig. 119. 



formed, and the latter, being volatile, must be condensed in a bottle 
surrounded with cold water. 

3C + 3C1 2 = 3C0 



APO 3 + 

Aluminium oxide. 



+ 2A1C1 3 

Aluminium chloride. 



Aluminium chloride is a white or pale-yellow crystalline com- 
pound, which melts at a gentle heat, and volatilizes at a tempera- 
ture slightly above 100°. When thrown into water, it dissolves 



ALUMINIUM SULPHATE. 303 

and combines with the liquid, forming a hydrate, which cannot be 
dried without decomposition. It slowly absorbs moisture from the 
air, giving off hydrochloric acid while aluminium oxide is formed. 
The vapor density of the chloride below 440° corresponds to 
the formula APC1 6 ; above that temperature it diminishes, and 
at 850° it is that required by the formula A1CP. 

508. Aluminium Oxide, APO 3 . — This compound is com- 
monly called alumina. It is found native in corundum, ruby, 
sapphire, oriental topaz, and emery : the black color of emery is due 
to the presence of oxide of iron. Aluminium oxide may be ob- 
tained in the laboratory by heating the aluminium hydrate which 
is thrown down as a gelatinous white precipitate when ammonia 
is added to a solution of alum. It then forms a white powder, 
which is infusible except in the oxyhydrogen flame and in the 
electric arc. It is not reducible by hydrogen, and carbon de- 
oxidizes it only in the electric furnace. The crystallized varie- 
ties of alumina are used as gems : ruby is red, sapphire is blue, 
and topaz is yellow. By reason of their hardness, corundum 
and emery are of great value in grinding and polishing glass, 
steel, and metals. 

Aluminium hydroxide, Al(OH) 3 , forms a bulky gelatinous pre- 
cipitate when ammonia-water or an alkaline hydroxide or carbonate 
is added to a solution of alum or any salt of aluminium. 

509. Aluminium Sulphate, AP(S0 4 ) 3 . — Clay is a silicate of 
aluminium, usually colored yellow by the presence of a little iron. 
By boiling with strong sulphuric acid, clay is decomposed, a solu- 
tion of aluminium sulphate being formed. This body is made 
from clay as free as possible from iron ; when its solution is evap- 
orated, a white crystalline mass is obtained, and by special precau- 
tions the salt may be crystallized in small pearly scales or needles 
containing eight molecules of water of crystallization. It is sol- 
uble in twice its weight of cold water, and is used as a mordant in 
dyeing, for it may be decomposed in the fibres of the tissues to be 
dyed, and the fine particles of aluminium oxide deposited firmly 
fix the color in the fabric. For this purpose it is usually first con- 
verted into aluminium acetate by the addition of calcium acetate. 



304 LESSONS IN CHEMISTRY. 

When aluminium sulphate is heated, it first loses itc water of 
crystallization, and then gives off sulphur trioxide, leaving a resi- 
due of aluminium oxide. 

A1 2 (S0 4 )» = 3S0 3 + A1 2 3 

510. Alums. — To a cold saturated solution of aluminium sul- 
phate, we add a cold saturated solution of potassium sulphate, and 
stir the mixture. A crystalline deposit forms. The two salts 
have combined to form a double salt, which is called an alum. It 
crystallizes with twenty-four molecules of water of crystallization, 
and its formula is A1 2 (S0 4 ) 3 .K 2 S0 4 + 24H 2 0. By the substitu- 
tion of sodium sulphate or ammonium sulphate for the potassium 
sulphate in the preceding experiment, sodium alum or ammonium 
alum will be formed. The compositions of these substances are 
precisely analogous to that of the potassium alum, and they crys- 
tallize in the same form, which is the regular octahedron. 

A12(S0 4 ) 3 .Na 2 S0 4 + 24H 2 Sodium alum. 
Al 2 (SO*) 3 .(NH*) 2 SO* + 24H20 Ammonium alum. 

Potassium alum is soluble in about thirty times its weight of cold water, and 
in less than one-third its weight of boiling water. It forms voluminous, trans- 
parent crystals when the hot saturated solution is allowed to cool. When 
heated, it melts in its water of crystallization, which is afterwards driven off; 
the salt increases enormously in volume, and the anhydrous alum then forms 
a white, porous mass. Alum may be obtained crystallized in cubes by adding 
a very small quantity of potassium carbonate or hydrate to its hot solution and 
allowing it to cool. 

Sodium alum is very soluble in cold water, and is not employed in the arts. 

Ammonium alum is the compound ordinarily called alum. Its solubility is 
about the same as that of potassium alum. When it is strongly heated, it 
leaves a residue of pure alumina. 

Other metals whose oxides are analogous in constitution to aluminium oxide, 

form alums having compositions and general properties like those of ordinary 

alum. These alums are isomorphous. Although their colors be different, they 

maybe mixed in the same crystal, and the form of the latter will remain 

unchanged. Thus, chromium alum is red : when one of its red octahedral 

crystals is immersed in a saturated solution of potassium alum and the water 

is allowed to evaporate, the octahedron will grow larger, and the red chromium 

alum will be surrounded by the colorless potassium alum. The compositions 

of three of these alums are shown by the following formulae : 

Iron alum, Fe 2 (S04) 3 .K 2 SO* + 24H 2 

Manganese alum, Mn 2 (SO*) 3 .K 2 SO* + 24H 2 

Chromium alum, Cr 2 (S0 4 ) 3 .K2SO* + 24H2Q 



CLAY AND POTTERY. 305 

511. Clay and Pottery. — The feldspars — orthoclase, albite } 
and labradorite — are double silicates of aluminium and potas- 
sium, sodium and calcium. The micas and garnets are also 
double silicates of aluminium and various other metals. The dis- 
integration of these minerals by the action of air and frost results 
in the formation of clays, and the nature of a clay will depend on 
that of the rock from which it is derived. The purest clay is a 
hydrated silicate of aluminium known as kaolin, or porcelain clay. 
It contains Al 2 3 .2Si0 2 .2H 2 0. Clays which form a coherent mass 
when mixed with water, and which when calcined become very 
hard without being fused, are called plastic clays, and are used for 
the manufacture of bricks, fire-brick, pottery, etc. Fuller's earth 
is a kind of clay of which the paste is not strongly coherent : it is 
used in scouring and fulling cloths. Marls are mixtures of clay 
and chalk, generally of a greenish color, and often found in large 
deposits : they are used as fertilizers for sandy soils. 

Porcelain is made from a mixture of the finest kaolin with a 
little finely-powdered sand and feldspar, which are added to pre- 
vent the mass from shrinking and to render the ware translucent 
by undergoing partial fusion. The greatest care is exercised that 
the materials, which are made into a paste with water, may be 
intimately mixed ; after the articles have been fashioned from the 
perfectly homogeneous paste, they are baked at a dull red heat, and, 
after cooling, are removed from the furnace. They are then dipped 
into water holding in suspension a mixture of kaolin and quartz in 
an impalpable powder. This powder fills the pores on the surface, 
and when the articles are again baked the mixture fuses and forms 
a transparent glaze, while the whole mass becomes partially vitrified. 

Stoneware is manufactured from a kaolin which is not sufficiently 
pure for porcelain-making. It is baked at one operation, and when 
the temperature of the oven is very high a little common salt is 
thrown on the incandescent objects; by the action of the hydrogen 
compounds in the flame, hydrochloric acid is formed, while the 
sodium forms a double silicate with aluminium on the surface of 
the ware. This silicate, being quite fusible, melts and spreads out 
on the surface of the ware, forming an even glaze. 

20 



306 LESSONS IN CHEMISTRY. 

Articles of faience are made from a still more common clay 
mixed with finely-powdered quartz, and, after being rendered cohe- 
rent by a preliminary baking, are coated with a mixture of pow- 
dered quartz, potassium carbonate, and lead oxide. This mixture 
fuses to a transparent varnish when the articles are baked a second 
time, and various colors are obtained by the addition of certain 
metallic oxides. Oxide of tin renders the glaze white and opaque. 

The glazing of pottery intended for culinary purposes should 
contain no lead, as lead silicate i& attacked by dilute vegetable 
acids, and a lead salt is sometimes so formed in articles of food. 

The blue pigment ultramarine is made by heating kaolin with sodium sul- 
phate, sodium carbonate, sulphur, charcoal, and rosin : the semi-vitrified mass 
is then pulverized, roasted, again pulverized, washed and dried. The mineral 
lapis lazuli is natural ultramarine. 

512. Tests for Aluminium. — Solutions of aluminium salts 
usually have an acid reaction. The alkaline hydroxides and am- 
monia produce gelatinous white precipitates of aluminium hydrox- 
ide, soluble in acids and in the alkaline hydroxides. The same 
precipitate is thrown down by the alkaline carbonates and by am- 
monium sulphide, carbon dioxide being liberated by the former 
and hydrogen sulphide by the latter. 
A1 2 (S0±) 3 + 3(NH±) 2 3 + 6H 2 = 2A1(0H) 3 + 3(NH 4 ) 2 S0± + 3H 2 S 

When an aluminium salt or aluminium oxide is strongly heated 
in a blow-pipe flame, and the resulting white mass is moistened 
with a drop of cobalt nitrate solution and again heated, it be- 
comes sky-blue, without fusion. 

513. Closely related to aluminium by chemical analogies are a number of 
jare metals, of which three have been obtained in the metallic state. They are 
cerium, didymium, and lanthanum. As silicates and phosphates they occur in 
cerite, fjadolinite, monazite, and other rare minerals : didymium has been 
resolved into two components, — neodymium and iiraseodymium. 

Scandium, samarium, hohnium, erbium, thulium, ytterbium, and. yttrium are 
elements which have not been obtained in the metallic state; but their oxides 
have been isolated in small quantities, and a few of their salts have been 
studied. Each of these elements is distinctly characterized by its spectrum, 
which is an unquestionable indication of the individuality of the element. 
All these elements appear to form sesquioxides, and their chlorides, which 
have been prepared, contain one atom of metal and three of chlorine, cor- 
responding in composition to aluminium chloride. 



IRON AND ITS METALLURGY. 307 



LESSON LVIII. 
IRON AND ITS METALLURGY. 

514. Iron is found in the metallic state in meteoric stones, which 
are occasionally drawn to the earth during its passage through space. 

The more important minerals from which iron is extracted 
are magnetite, Fe 3 4 ; hematite, or specular iron, Fe 2 3 ; Union .ite, 
or brown hematite, 2Fe 2 3 -f 3H 2 ; goethite, Fe 2 3 + H 2 ; 
and siderite, or spathic iron, FeCO 3 . Iron pyrites, which is 
chiefly employed for the manufacture of sulphuric acid, is the 
disulphide FeS 2 . Some of the minerals are found in a tolerably 
pure condition, but generally they are mixed with silicious mat- 
ters, clay, limestone, coal, etc. When the ore contains sulphur, 
it is first roasted, and the sulphur is burned into sulphur dioxide, 
while the iron remains as oxide. Very frequently also the ores 
are calcined to expel water and carbon dioxide, and to render 
them more porous. 

The oxide of iron, either the natural ore or produced by the 
roasting, is reduced by being heated with charcoal. A rather primi- 
tive method, but one which for ages furnished all of the iron em- 
ployed, and which is still used for the reduction of very rich ores, 
is known as the Catalan' method, the name being derived from 
the Spanish province in which the process is still carried on. It 
consists in piling the ore and charcoal in two heaps, side by side, 
on burning charcoal contained in the hearth of a furnace where 
the combustion is sustained by a blast of air from the tuyere of a 
bellows (Fig. 120). The reduced iron collects on the hearth in a 
spongy mass, which is removed and directly submitted to the 
operation of forging. The silicious matters of the ore combine 
with a portion of ferrous oxide produced during the operation, and 
form a very fusible slag, consisting of ferrous silicate. 

515. The blast-furnace process for the reduction of iron is 



308 



LESSONS IN CHEMISTRY. 



applicable to all iron ores, and the fuel employed is either char- 
coal, coke, or anthracite. The blast-furnace is a tall column of 
considerable height, sometimes almost cylindrical, but more often 

constructed in the form of a 
double frustum of cones placed 
base to base (Fig. 121). It is 
lined with infusible fire-brick ; 
the hearth is flat, and inclines 
very slightly towards the front, 
which is so arranged that the 
molten iron may be drawn off 
at the bottom through a hole 
which is kept closed with a clay 
plug, and the slag may be re- 
moved as it accumulates and 
floats on the surface of the liquid 
iron. During the reduction, the 
bottom of the furnace is closed, 
and a blast of air is injected 
through tuyere pipes (T) by 
powerful blowing engines. Coal, 
ore, and limestone are continually supplied in alternate layers at the 
open top of the furnace, and in the interval between the introduc- 
tion of these materials the top is closed by a conical cap or dome, 
which can be readily moved by suitable machinery. At the be- 
ginning of the operation the furnace is heated by a supply of com- 
bustibles only, and, when the temperature has been sufficiently 
raised, the ore and limestone are gradually alternated with the 
introduction of coal, until the furnace is completely filled. By 
the combustion of the coal immediately above the tuyeres, carbon 
dioxide is produced, but as this comes in contact with the highly- 
heated coal it is reduced to carbon monoxide : as the latter gas 
rises through the mixture, it reduces the oxide of iron, and the 
metallic iron formed is disseminated in small particles through the 
mass of reduced ore. As the materials descend in the furnace, 
the silicious matters of the ore, and the lime resulting from the 




Fig. 120. 



BLAST-FURNACE PROCESS. 



309 



action of the heat on the limestone, unite to form a very fusible 
slag of calcium and aluminium silicate, while the particles of iron 
are agglomerated together, and, together with the slag, flow to the 
hearth of the furnace. The gases produced during the operation 
contain a large proportion of carbon monoxide ; they are carried 




Pig. 121. 

off by pipes inserted near the top of the furnace, and their com- 
bustion furnishes heat for the boilers which supply steam to the 
blowing engines and to furnaces containing long series of pipes, 
through which the air from the engines is forced before it enters 
the tuyeres. The blast is heated as highly as possible, for by the 
use of hot air a very great saving is effected in the quantity of 
fuel required. 



310 LESSONS IN CHEMISTRY. 

When a sufficient quantity of iron has accumulated on the 
hearth of the furnace, the blowing engines are slowed or stopped, 
and by picking out the clay plug the molten iron is caused to flow 
into semi-cylindrical channels in sand on the floor of the casting- 
room. Blows from a sledge-hammer detach the bars of iron so 
formed from that in the channel from which the moulds are filled, 
and they constitute pig-iron. 

This iron contains carbon and smaller proportions of silicon, 
sulphur, and phosphorus, which are derived from the various 
materials in the blast-furnace. These impurities are in great part 
removed by melting the iron in puddling-furnaces, where the 
carbon and silicon are oxidized either by air, or better by oxygen 
derived from pure magnetic iron ore, or from scales of black oxide 
of iron obtained in another operation. During the process, the 
molten iron is vigorously stirred until it is converted into a spongy 
mass, which is then removed and placed under a steam-hammer, 
by the blows of which all the ferrous silicate and black oxide of 
iron formed during the process of puddling are squeezed out, while 
a bloom of soft iron remains. The ferruginous scoriae or ashes 
obtained in this operation are used in refining a new quantity of 
cast-iron. 

516. The soft iron of commerce, known as wrought-ivon, is 
not perfectly pure. It contains traces of carbon, silicon, sulphur, 
phosphorus, nitrogen, and sometimes other elements. Pure iron 
may be obtained by passing hydrogen over pure ferric chloride 
heated to bright redness in a porcelain tube. Hydrochloric 
acid is disengaged, and the iron remains as an almost infusible, 
spongy mass. By passing dry hydrogen over ferric oxide heated 
to dull redness in a glass bulb (Fig. 122), metallic iron is ob- 
tained as a dull black powder, in which form it becomes oxidized 
with great readiness. If a lighted match be applied to a single 
point in recently-prepared iron reduced by hydrogen, the whole 
mass quickly takes fire and burns into ferric oxide. By reducing 
the ferric oxide at a temperature below redness, a powder of 
iron may be obtained which will even take fire spontaneously on 
contact with the air. 



STEEL. 



311 



The composition and general properties of cast-iron vary 
greatly, for while cast-iron always contains silicon and carbon, 
these elements are only in part chemically combined with 
the iron. The proportion of carbon varies from 2 to 5.5 per cent. 
When cast-iron containing a large proportion of carbon is rapidly 
cooled, it becomes hard and brittle, and its fracture is coarse, crys- 
talline, and very white. It is called white iron. When, however, 
such iron is allowed to cool slowly, a considerable quantity of the 




Fig. 122. 

carbon separates as shining scales of graphite, and the iron is then 
softer, has a closer structure, and a gray fracture. It is called 
gray iron. Iron containing sulphur and phosphorus is always 
white ; phosphorus renders iron brittle while cold, and sulphur 
renders it brittle while hot. In the first case the iron is said to 
be cold-short, while in the second it is called red-short. 

Spiegeleisen and ferro-manganese are varieties of cast-iron 
very rich in carbon, and containing much manganese. They 
are employed in the manufacture of steel. Spiegel is crystal- 
line, and breaks with smooth and highly-lustrous surfaces. 

517. Steel is iron containing from 0.2 to 2 per cent, of carbon, 
and traces of nitrogen. It is obtained by a number of processes, 
which depend either on the partial decarbonization of cast-iron or 
on the introduction of the required proportion of carbon into soft 



312 



LESSONS IN CHEMISTRY. 



iron. When manganiferous cast-iron is maintained for a time 
melted under a layer of magnetic iron ore or ferruginous scoriae, 
and the operation is arrested at the proper moment, the iron still 
retains a certain proportion of carbon, and natural steel is obtained. 
Cement-steel, or blister-steel, is made by piling soft-iron bars 
between layers of charcoal in fire-clay boxes, which are then 
heated to redness in a furnace, and the temperature is maintained 
for several days ; the iron absorbs a certain proportion of carbon, 
and is converted into steel. As, however, the exterior of the bars 
will necessarily contain more carbon than the interior, the metal 
is rendered homogeneous by being melted in crucibles heated in a 
powerful wind-furnace. It then constitutes cast-steel. 

The most important method of manufacture of steel is named, 
from its inventor, the Bessemer process. It consists in completely 
decarbonizing cast-iron and then adding sufficient cast-iron of the 
proper quality to give to the whole mixture the desired proportion 
of carbon. The operation is conducted in oval vessels of strong 

iron plate lined with infusi- 
ble fire-brick. This appa- 
ratus, which is called a 
converter, is supported on 
trunnions, so that it may 
swing back and forward on 
a horizontal axis. One of 
the trunnions is hollow, and 
communicates with a pipe 
passing partly around the 
converter and then leading 
to its bottom ; the fire-brick 
is here pierced with a 
number of holes, so that a 
HP blast of air may be forced 



Fig 123 U P ^ nrou § n ^ De contents of 

the converter (Fig. 123). 

The plant, as the whole of any manufactory is called, is established 

near blast-furnaces, and cupola furnaces, for melting the pig-iron, 




BESSEMER STEEL PROCESS. 313 

are constructed near the converter. Everything being ready, 
burning wood is thrown into the converter, which is then partly 
filled with coke, and the blast is turned on so that the fire-brick 
lining is heated to whiteness. The converter is then inverted ; 
the coke is dumped out, and molten iron is run in from cupola 
furnaces. During the filling, the converter is kept in an inclined 
position, so that the tuyeres for the passage of the blast may not 
become filled with the molten iron. The blast, which is under 
strong pressure, is now turned on, .and the converter is rotated 
to an upright position. As the air bubbles up through the iron, 
the carbon, silicon, and other oxidizable elements present are con- 
sumed, and a brilliant flame rushes with a roaring noise from the 
mouth of the converter. When the carbon of the iron is burned 
out, the appearance of the flame undergoes a change which informs 
the workmen of the termination of the operation. The converter 
is then inclined, and the blast is arrested. In the mean time, the 
quantity of iron in the charge being accurately known, the exact 
quantity of spiegel required to convert the charge into steel" has 
been melted in another cupola furnace : the blast is stopped, the 
molten spiegel run into the converter, and the blast turned on 
for a moment in order that the contents maybe perfectly mixed. 
The steel is then poured out into an enormous ladle, which is 
carried by a revolving crane over the circumference of a circle 
around which are arranged ingot-moulds, into which the steel is 
cast. 

A modification of the Bessemer process, known as the basic pro- 
cess, is extensively used in Europe for making steel from pig- 
iron rich in phosphorus. The fire-brick lining of the converter 
is replaced by one of dolomite, the bases of which aid in the oxi- 
dation and removal of the phosphorus by forming phosphates. 
These constitute a slag valued as a fertilizer in agriculture. 

518. The valuable qualities of steel depend upon the ease with 
which it can be hardened or softened at pleasure, and the opera- 
tions bv which the change in hardness is brought about consti- 
tute the processes of tendering. When heated and allowed to 
cool slowly, steel becomes soft and malleable like soft iron, but if 



314 LESSONS IN CHEMISTRY. 

it be heated to redness and then suddenly cooled by plunging it 
into cold water, it is rendered hard and brittle ; it is, however, 
still elastic. Intermediate degrees of hardness are obtained by 
re-heating to temperatures depending on the desired hardness. 
Part of the temper is then said to be drawn. The color which 
the surface of the metal assumes is an index of the temperature. 

Straw yellow corresponds to 230°. 
Brown " 255°. 

Light blue " 285-290°. 

Indigo blue " 295°. 

Sea green " 331°. 

When a number of small objects are to be tempered alike, they 
are heated in a bath of mercury or oil, of which the temperature 
earn be exactly regulated. 

We may very well study the phenomena of tempering by heat- 
ing a steel wire or a piece of watch-spring to redness and quickly 
immersing it in water. It is now very hard and brittle ; it breaks 
as readily as a piece of glass. We again heat it gently until its 
surface becomes of a blue color, and now, whether we dip it in 
water or allow it to cool slowly, we will find that it has become 
quite elastic, but is still hard. When, however, we heat it to red- 
ness and allow it to cool slowly, it becomes soft and flexible ; it 
will not break, but will retain any form into which it is bent. 

The process of casehardening, which is applied to inferior 
kinds of cutlery, consists in embedding the objects, which are 
made of soft iron, in charcoal contained in crucibles or clay boxes. 
These are then heated to bright redness, and the surface of the 
iron becomes converted into steel. 

Steel is less fusible than cast-iron, but much more fusible than 
soft iron ; at the temperature at which soft iron becomes pasty, 
steel melts. 



IRON AND ITS COMPOUNDS. 315 

LESSON LIX. 
IRON AND ITS COMPOUNDS. 

519. The density of soft iron varies from 7.4 to 7.9. It is 
ductile, malleable, and very tenacious. It fuses only at the highest 
temperatures of a powerful wind-furnace, but at a high white 
heat it becomes so soft that two pieces of the metal may be readily 
united in a solid mass by hammering or by pressure : the opera- 
tion is called welding. At a somewhat lower temperature it may 
be readily rolled into sheets or bars, and sheet-iron is made by 
passing the heated metal between polished steel rollers. Iron may 
be rolled into leaves as thin as paper. Tin plate is sheet-iron 
covered with a coating of tin. Galvanized iron is made by dipping 
perfectly clean sheet-iron into melted zinc. 

Iron is attracted by a magnet, and becomes itself a magnet 
while under the magnetic influence, but loses its magnetism when 
the exciting cause is removed. Under the same circumstances 
steel becomes a permanent magnet. 

Unless in a state of fine division, iron is unaffected by dry air 
at temperatures below redness, but at a red heat it combines with 
oxygen and is converted into a black oxide which forms scales on 
its surface. It is rapidly rusted by moist air, and the rust is a 
hydrated ferric oxide. 

When the formation of rust has begun, it proceeds with great rapidity, and 
if the mass of iron be of such a form that a large surface is combined with a 
comparatively small bulk, as in a long coil of wire or mass of small scrap iron 
partly immersed in water, the temperature may be much elevated by the rust- 
ing. It appears that hydrogen dioxide is formed during the rusting of iron, and 
that substance would greatly accelerate the change : the nitrogen of the air 
also plays some part in the phenomenon, for rust always contains a trace of 
ammonia. 

At a red heat, iron decomposes water, liberating hydrogen, and 

forming an oxide. It is dissolved by hydrochloric and sulphuric 

acids, hydrogen being set free j this hydrogen has an unpleasant 



316 LESSONS IN CHEMISTRY. 

odor, probably due to carbon compounds formed by the action 
of the carbon of the iron. Dilute nitric acid also dissolves iron, 
disengaging red vapors, but the strongest nitric acid does not 
affect it. 

If some clean iron wire or some bright nails be dropped into pure nitric acid, 
or a mixture of strong nitric and sulphuric acids, no action takes place; the 
iron may now be removed and placed in more dilute acid, and even here it 
will not dissolve : it is said to be in the passive state. Its surface has become 
covered with a protecting layer of gas derived from the strong acid : if while 
the passive iron is immersed in the dilute acid we touch its surface with a 
copper wire, the coating of gas is broken at one point, chemical action is at 
once re-established, and the iron is quickly dissolved. 

Iron forms two series of compounds, — ferric compounds, in 
which the iron is triatomic, and ferrous compounds, in which it 
plays the part of a diatomic element. Under certain conditions 
two atoms of iron act as a hexatomic couple. 

520. Ferrous Chloride, FeCl 2 , is made by passing dry hydro- 
chloric acid gas over metallic iron heated to redness in a porcelain 
tube. It then condenses in white, pearly scales in the cooler part 
of the tube. 

A solution of ferrous chloride may be obtained by dissolving iron in hydro- 
chloric acid. When the filtered liquid is sufficiently evaporated, it deposits 
bluish-green crystals in which every molecule of ferrous chloride is combined 
with four molecules of water. 

521. Ferric Chloride, FeCP, sublimes in brilliant violet 
crystals when chlorine is passed over incandescent iron con- 
tained in a glass or porcelain tube. It is very soluble in water, 
but its solution undergoes a curious change by boiling. A solu- 
tion of ferric chloride is obtained by dissolving ferric oxide in 
hot hydrochloric acid. When this solution is evaporated at a 
low temperature, the hydrated ferric chloride remains as a 
brownish-yellow, deliquescent mass, but when the solution is 
boiled its color darkens, and the reactions and general properties 
of the liquid seem to show that it has been decomposed into 
hydrochloric acid and a soluble variety of ferric hydroxide. 
Above 700° the vapor density of ferric chloride corresponds to 
the formula FeCP, but at 450° it is that required by Fe 2 Cl 6 . 



OXIDES OF IRON. 317 

522. Ferrous Oxide, FeO, has been obtained as a black powder by passing 
a mixture of carbon dioxide and carbon monoxide in equal volumes over heated 
ferric oxide. Carbon monoxide alone would yield metallic iron. 

523. Ferric Oxide, Fe 2 3 , constitutes the minerals known as 
red hematite and specular iron. It is obtained as a fine red powder 
by strongly heating ferrous sulphate in a crucible : sulphur dioxide 
and sulphur trioxide are disengaged, while ferric oxide remains. 

2FeSO* = SO 2 + SO 3 + Fe 2 3 
This powder is very hard, and is used for polishing under the 
names jewellers' rouge and colcothar. 

When an alkaline hydroxide or ammonia is added to a solution 
of ferric chloride, a flocculent, brown precipitate of ferric hydrox- 
ide, Fe(OH) 3 , is thrown down. This is the precipitate which, 
after being thoroughly washed, is the proper antidote for poison- 
ing by arsenious oxide. Ferric solutions containing tartaric acid 
are not precipitated by the alkaline hydroxides. 

Rust is a ferric hydroxide of which the composition usually corresponds with 
the formula (Fe 2 3 ) 2 3H 2 = 2Fe(OH) 3 + Fe 2 3 . This is also the composition 
of the natural hydrate brown hematite. Goethite is a hydroxide having the 
composition FeW.H 2 = 2FeO(OH). 

There is a soluble modification of ferric hydroxide. It may be obtained by 
pouring a solution of ferric chloride which has been heated to 100° into the 
inner vessel of a dialyser ($ 220), the water in the exterior vessel being fre- 
quently changed. Hydrochloric acid passes through the membrane, while a 
solution of ferric hydroxide remains within. Dialysis of a solution of ferric 
acetate yields soluble ferric hydroxide in the same manner. This solution is 
used in medicine under the name dialysed iron. 

524. Ferroso ferric Oxide, Fe 3 0\ is magnetic oxide of iron, 
commonly called black oxide of iron. It is found native in large 
quantities in the neighborhood of Lake Superior. It forms in 
black scales on the surface of iron heated to redness in the air. It 
is attracted by the magnet. It is a compound of ferrous and ferric 
oxides, Fe 3 4 = FeO.Fe 2 3 . 

525. Ferrous Sulphide, FeS, so largely used in the labora- 
tory for the preparation of hydrogen sulphide, is made by heating 
a mixture of iron filings with two-thirds its weight of sulphur to 
redness in a covered crucible. After fusion, the mass is poured 
out, and on cooling forms a black solid of a metallic appearance. 

526. Iron Disulphide, FeS 2 , constitutes the common mineral 



318 LESSONS IN CHEMISTRY. 

iron pyrites. It is dimorphous, being found in cubical crystals 
of a yellow color and metallic lustre, known as pyrite ; and as 
rhombic prisms of a pale, greenish-yellow color, constituting mar- 
casite. When pyrites is heated in closed vessels, part of its sul- 
phur distils ; when it is heated in contact with air, the sulphur 
burns into sulphur dioxide, while the iron remains as oxide. The 
brilliant metallic appearance of iron pyrites has sometimes caused 
it to be mistaken for gold, and it has been called fool's gold : the 
action of heat at once reveals its true character. 

527. Ferric Sulphate, Fe 2 (S0 4 ) 3 . — Ferrous sulphate, or 
green vitriol, has already been described (§ 123) : when crystals 
of this salt are dissolved in water, and boiled with a little less than 
one-sixth their weight of sulphuric acid, and small quantities of 
nitric acid are added from time to time, a solution of ferric sul- 
phate is obtained. When this liquid is evaporated to dryness, 
ferric sulphate remains as a yellowish-white mass, very soluble in 
water. By using a smaller quantity of sulphuric acid, various 
basic salts are obtained, and they may be considered as ferric sul- 
phate in which one or two groups, SO 4 , are replaced by as many 
atoms of oxygen. Such are Fe 2 0(S0 4 ) 2 and Fe 2 2 SO. A mix- 
ture of these basic sulphates is employed in medicine under the 
name Monsel's solution. It is astringent and styptic, and is valu- 
able for arresting hemorrhage. 

528. Tests for Iron. — The ferrous and the ferric salts are 
characterized by different reactions ; by reducing agents such as 
nascent hydrogen produced by zinc and hydrochloric acid, the ferric 
salts are converted into ferrous salts, while ebullition with nitric 
acid, or the addition of chlorine- water, will produce a ferric com- 
pound from a ferrous salt. 

Solutions of the ferrous salts are pale green ; hydrogen sulphide 
occasions in them no precipitate, but ammonium sulphide throws 
down black ferrous sulphide. The alkaline hydroxides and ammonia 
produce greenish- white precipitates of ferrous hydroxide which rap- 
idly become dark by absorbing oxygen from the air. Potassium 
ferrocyanide forms a white precipitate instantly changing to pale- 
blue ; potassium ferricyanide a dark-blue precipitate, called Turn- 
bull's blue. 



COBALT. 319 

Solutions of ferric salts are yellowish-brown or brown. With 
hydrogen sulphide they yield a precipitate of sulphur, being 
reduced to ferrous salts ; ammonium sulphide throws down a 
black precipitate. The alkaline hydroxides and ammonia form rust- 
colored precipitates of ferric hydroxide, insoluble in an excess of the 
reagent. Potassium ferrocyanide throws down Prussian blue ; 
potassium ferricyanide occasions no precipitate. Potassium sul- 
phocyanate produces a blood-red color, due to the formation of 
ferric sulphocyanate. Tannin, or an infusion of gall-nuts, forms 
a blue-black and very finely divided precipitate, which long remains 
suspended in the liquid. 



LESSON LX. 



COBALT, NICKEL, AND MANGANESE. 

529. Cobalt, Co = 59. — This metal is found combined with 
sulphur and arsenic ; smaltite, CoAs 2 , and cobaltite, CoAsS, are 
its principal ores. Much cobalt is also extracted from earthy 
cobalt, or asbolite, a mixture or compound of oxide of cobalt 
with hydrated oxides of manganese. The metal is obtained by 
reducing the oxides with carbon or hydrogen. 

Pure cobalt is a grayish- white metal, having a reddish tinge 
and a strong lustre. It is malleable and ductile. It melts at 
about 1800°, and has a density of 8.6. It is attracted by the 
magnet. Neither dry nor moist air affect it at ordinary tem- 
peratures, but it is oxidized at a red heat. 

Cobalt forms cobaltous oxide, CoO, a sesquioxide, Co 2 3 , and several other 
oxides which appear to be formed by a combination of these two in different 
proportions. It is in the form of these oxides that cobalt is separated from 
its ores. 

530. Cobalt Chloride, CoCl 2 , is prepared by dissolving either the oxide or 
carbonate of cobalt in hydrochloric acid. The solution is red, and, when con- 
centrated, deposits red crystals containing six molecules of water of crystalli- 
zation. Anhydrous cobalt chloride is blue : if a little strong sulphuric or 
hydrochloric acid be added to a concentrated solution of cobalt chloride, the 



320 LESSONS IN CHEMISTRY. 

liquid becomes blue. It contains anhydrous cobalt chloride. Writing made 
on paper with a very dilute solution of cobalt chloride is invisible when dry ; 
the small quantity of the salt present is hydrated ; but if the paper be heated, 
the characters become blue, for the water is driven off. After exposure to the 
air for a time, the characters again fade, the cobalt chloride absorbing atmos- 
pheric moisture. The solution is employed as a sympathetic ink. 

531. Cobalt Blue. — The ores of cobalt are principally em- 
ployed for the manufacture of a dark-blue substance generally 
called smalt. This is a mixture of cobalt silicate and potassium 
silicate. It is prepared by partially roasting the ore in order to 
convert the greater part of the cobalt into oxide. The roasted 
mass is then pulverized and melted with a mixture of potassium 
carbonate and white quartz sand. A blue, vitreous mass is thus 
obtained, which floats on a fused mass containing the iron, nickel, 
copper, and unaltered sulphur and arsenic of the ore. This mix- 
ture has a metallic appearance ; it is called speiss, and is used for 
the preparation of nickel. While still molten, the blue glass con- 
stituting smalt is poured into water, in which it breaks up into 
small fragments, which are readily pulverized. 

532. Tests for Cobalt. — The more ordinary salts of cobalt form rose- 
colored or currant-red solutions, but if these solutions contain free acid, they 
become blue when heated. They are not precipitated by hydrogen sulphide 
in acid liquids, but in neutral solutions black CoS is thrown down ; ammonium 
sulphide forms a black precipitate. The alkaline hydroxides produce blue pre- 
cipitates, which by boiling are converted into rose-colored cobaltous hydroxide, 
Co(OH) 2 . Ammonia-water occasions a blue precipitate, which dissolves in an 
excess of the reagent, an ammonio-cobalt salt being formed. When strongly 
heated with a little borax on the end of a platinum wire, the compounds of 
cobalt yield beads of a blue glass. 

533. Nickel, Ni — 59. — Nickel is found as arsenide in nicco- 
lite or kupfernickel, NiAs, as sulpharsenide in gersdorfflte, NiAsS, 
as sulphide in millerite, NiS, and in nickeliferous pyrites, and 
as a hydrated silicate in garnierite. The metal is extracted 
by various processes from these minerals and from the speiss 
formed during the manufacture of smalt. 

Pyrites, containing small proportions of nickel and copper, are treated as 
described under Copper (see p. 284), and so converted into a matte in which 
the two metals are concentrated, while most of the iron passes into the slag. 
Arsenical ores and speiss are first oxidized by roasting and then fused (smelted) 
with sand and soda. The products are a speiss, consisting of the arsenides of 
nickel, cobalt, and copper, and a slag which contains the iron and some of the 



NICKEL. 321 

cobalt in the form of silicates and arsenates. The whole process is then re- 
peated with the speiss to increase the proportion of nickel. 

The separation of nickel from matte and speiss is effected by either a dry- 
process or a wet process. In the former sulphur and arsenic are removed by 
roasting, and fusion with soda and saltpetre, nickel oxide remaining after 
lixiviating the fused mass with water. A better wa}^ is to roast to thorough 
oxidation, and to dissolve the resulting oxides in sulphuric acid. From the 
solution the copper is removed by means of hydrogen sulphide ; the iron is con- 
verted into the ferric state and precipitated by calcium carbonate, and cobalt 
thrown down as sesquioxide by adding bleaching powder to the neutral solution. 
Upon addition of milk of lime to the decanted liquid, the nickel is obtained 
as the hydroxide, Ni(OH) 2 , which, by ignition, is converted into the oxide. 

The oxide, after being purified, is reduced by mixing it with charcoal and 
heating the mixture to a high temperature. The resulting metal always con- 
tains carbon and small amounts of metallic impurities. 

Pure nickel may be obtained by strongly heating the oxalate 
out of contact with the air, by reducing the pure oxide by 
hydrogen, by electrolvzing solutions of nickel salts, and by de- 
composing nickel carbonyl, Ni(CO) 4 (see p. 153), by heat. 

Nickel is a silver- white metal, capable of taking a high polish. 
It is malleable, ductile, and very tenacious. Its density is about 8.5. 
It is attracted by the magnet. It is the hardest of the more common 
metals. It is not affected by the air at ordinary temperatures, but 
becomes oxidized at a red heat. It is slowly dissolved by dilute 
hydrochloric and sulphuric acids, more rapidly by nitric acid. 

Nickel is employed in the manufacture of a number of alloys. 
German silver contains 25 per cent, of nickel, 25 per cent, of 
zinc, and 50 per cent, of copper, but the proportions vary 
greatly. The white nickel coins of the United States contain 
25 per cent, of nickel and 75 per cent, of copper. Nickel-steel, 
an alloy of nickel and iron, is extensively used for armor-plates, 
and in the construction of heavy machinery. 

Nickel is largely employed for plating articles of brass, iron, 
and steel, and its hardness, its high lustre, and its freedom from 
rust render it admirably adapted to this purpose. The well- 
cleaned objects are attached to the zinc pole of a voltaic battery 
and immersed in a solution of nickel and ammonium double 
sulphate : the positive pole of the battery is connected with a 
plate of pure nickel dipped in the same liquid. 

21 



322 LESSONS IN CHEMISTRY. 

534. Nickel Chloride, NiCl 2 , is made by dissolving the oxide or hydrate 
in hydrochloric acid. When sufficiently concentrated, the green solution 
deposits green crystals containing NiCl 2 + 6H 2 0. 

535. Nickel Monoxide, NiO, is a green powder, obtained by strongly heat- 
ing the carbonate or nitrate. Nickel hydroxide, Ni(OH) 2 , is thrown down 
as a pale-green precipitate when an alkaline hydrate is added to the solution 
of a nickel salt. When chlorine is passed through water in which this pre- 
cipitate is suspended, a hydrate of nickel sesquioxide, Ni 2 3 , is formed. 

536. Nickel Sulphate, NiSO 4 . — When nickel oxide or hydroxide is dissolved 
in dilute sulphuric acid, and the solution is allowed to evaporate spontaneously, 
green crystals of the sulphate with seven molecules of water of crystallization 
are deposited. With ammonium sulphate, this compound forms a double salt 
in fine bluish-green crystals containing NiS0 4 .(NH 4 ) 2 S0 4 + 6H 2 0. This is 
the salt used in nickel-plating. 

537. Tests for Nickel. — The anhydrous nickel salts are yellow, but the 
crystallized salts and their solutions are emerald-green. If the solution be 
acid, hydrogen sulphide produces no precipitate, but a black precipitate of 
sulphide is thrown down if the solution contain sodium acetate. The same 
precipitate is formed by ammonium sulphide. The alkaline hydroxides and 
carbonates occasion pale-green precipitates. Ammonia-water forms a green 
precipitate, which dissolves in an excess of the reagent, yielding a blue solution. 

538 Manganese, Mn = 55. — This metal has been obtained 
by reducing the oxides with either carbon or aluminium at high 
temperatures. Its properties are greatly modified by small pro- 
portions of impurities. Pure manganese appears to be soft and 
malleable ; the presence of carbon renders it hard and brittle. 
The metal has a steel-gray color ; it melts and volatilizes in the 
electric arc. It is attacked by moist air, especially when it 
contains carbon. Three of its oxides are found native in the 
minerals braunite, Mn 2 3 , pyrolusite, MnO 2 , and hausmannite, 
Mn 3 4 . 

539. Manganese Dioxide, MnO 2 , is commonly called black 
oxide of manganese. When it is heated to redness, it loses oxy- 
gen, and is converted into red manganeso-manganic oxide, Mn 3 Q 4 . 
3Mn0 2 = Mn 3 4 + O 2 

Oxygen is also evolved, while manganous sulphate is formed, 

when the dioxide is heated with sulphuric acid. 

2MnO* + 2H 2 S0 4 = 2MnS0 4 + 2H 2 + 0' 
When heated with hydrochloric acid, manganese dioxide yields 

manganese chloride, water, and chlorine : large quantities of the 



MANGANIC ACID. 32:> 

dioxide are used for the manufacture of chlorine, and in the 
solution of manganese chloride the dioxide is regenerated. The 
process consists in mixing the liquid with milk of lime, by which 
calcium chloride and manganous hydroxide, Mn(OH) 2 , are 
formed : heated air is then blown through the mixture, and the 
manganous hydroxide is converted into the dioxide, from which 
the solution of calcium chloride is decanted. 

Manganese dioxide is also employed to decolorize glass ren- 
dered dark by carbonaceous matter or green by iron. In the 
first case it oxidizes the carbon, and in the second it converts 
the ferrous into ferric silicate, of which the yellow tint is neu- 
tralized by the purple color of manganic silicate. 

540. Manganic Acid, H 2 Mn0 4 . — When strongly heated with 
alkaline hydroxides, manganese dioxide absorbs oxygen from the air, 
and an alkaline manganate is formed. A mixture of manganese 
dioxide and potassium hydroxide may be fused in a silver or iron 
dish, and, when the cold mass is treated with water, a green solu- 
tion is obtained. If this be evaporated at a low temperature in a 
vacuum, it deposits green crystals of potassium manganate. 

When the alkaline manganates are heated to 450° in a current 
of steam, they are decomposed into alkaline hydroxide and manga- 
nese dioxide. This decomposition has been applied to the manu- 
facture of oxygen on a large scale. A mixture of sodium hydroxide 
and manganese dioxide is heated in a current of air ; oxygen is 
absorbed, and sodium manganate is formed. 

MnO 2 + 2NaOH + = Na 2 Mn0 4 + H 2 
Manganese dioxide. Sodium manganate. 

The air is then stopped, and steam is passed over the heated 

manganate, reproducing sodium hydroxide and manganese dioxide, 

while oxygen is disengaged. 

Na 2 Mn0 4 + H 2 = 2NaOH + MnO 2 + 
The oxygen, with the excess of steam, is led through cold 

pipes, where the steam is condensed, while the oxygen passes on 

to appropriate gas-holders. 

541. Permanganates. — When the green solution of potas- 
sium manganate is boiled, its color changes to red, while hydrated 



^24 LESSONS IN CHEMISTRY. 

manganese dioxide separates in brown flakes. The red color is 
due to the formation of potassium permanganate, and the solu- 
tion contains free potassium hydroxide. 

3K»MnO* + 2H20 = 2KMn0 4 + MnO 2 + 4KOH 

Potassium manganate. Potassium permanganate. 

A similar reaction takes place when an acid is added to the 
solution of a manganate. 

Potassium permanganate is made by heating in an iron crucible 
a mixture of five parts of potassium hydroxide with a little water, 
and three and a half parts of potassium chlorate with four of 
manganese dioxide in fine powder. The temperature is gradually 
raised to dull redness, the mass being constantly stirred. It is 
allowed to cool, and, after being pulverized, is thrown into two 
hundred parts of boiling water, and stirred until the liquid has 
assumed a purple color. It is then left to settle ; the clear liquid 
is decanted, neutralized with nitric acid, and evaporated on a 
water-bath. The crystals which separate on cooling are drained 
on a clean brick. 

They are purple-black needles, having a metallic reflection, 
soluble in about fifteen times their weight of cold water. The 
solution has an intense purple color. Potassium permanganate is 
an energetic oxidizing agent ; its solution is at once decolorized by 
sulphur dioxide, which it converts into sulphuric acid, and the 
liquid contains sulphuric acid, potassium sulphate, and manganous 
sulphate. 

2KMnO + 5S0 2 + 2H 2 = K 2 S0 4 + 2MnS0 4 + 2H 2 S0 4 

Potassium permanganate. Manganous sulphate. 

The oxidizing properties of potassium permanganate are largely 
employed in the laboratory. 

542. Tests for Manganese. — The salts of manganese are 
colorless or pale rose-colored. They are not precipitated by hydro- 
gen sulphide, but ammonium sulphide throws down flesh-colored 
manganese sulphide. The alkaline hydroxides produce dirty- white 
precipitates of manganese hydroxide, which soon absorbs oxygen 
from the air and becomes brown When heated with a little 
potassium hydroxide or nitrate or sodium carbonate on a piece of 



CHROMIUM. 325 

platinum foil, they yield a bluish- green mass of an alkaline 
manganate, which forms a red solution when treated with a little 
dilute nitric acid. 



LESSON LXL 

CHROMIUM AND TIN. 

543. Chromium, Cr = 52.5. — Chromium exists in the min- 
eral chromite, or chrome iron, which is a compound of chromium 
oxide and ferrous oxide, and may be considered as ferroso-ferric 
oxide in which the ferric oxide is replaced by the sesquioxide 
of chromium, FeO.Cr 2 3 . It is found also in the mineral cro- 
coite, PbCrO 4 . The metal is obtained by reducing the oxide 
either by means of carbon in the electric furnace, or by heating 
it with aluminium. Chromium is a grayish- white, very brilliant 
metal, having a density of 6.92. It is rather soft, but a small 
proportion of carbon renders it exceedingly hard. The pure 
metal is not magnetic, and very difficult to fuse. 

544. Chromic Chloride, CrCl 3 , is obtained by passing chlorine gas over 
an incandescent mixture of chromium sesquioxide and charcoal ; it then sub- 
limes and condenses in brilliant violet scales in the cooler parts of the tube. 
By the action of hydrogen at a red heat, it is converted into white chromous 
chloride, CrCl 2 . Chromic chloride is insoluble in water, but dissolves readily 
in presence of a small quantity of chromous chloride, yielding a green solution 
from which there may be obtained a crystallized hydrate, CrCl 3 -f 6H 2 0. 

545. Chromium Sesquioxide, Cr 2 3 . — When potassium dichromate is heated 
in a crucible with about half its weight of sulphur, a mass is obtained from 
which water dissolves potassium sulphate, leaving chromium sesquioxide as a 
green powder. Instead of sulphur, starch in quantity equal to one-fourth 
the weight of the dichromate may be employed, but the resulting oxide must 
afterwards be recalcined in the air, to burn out traces of carbon. It is in the 
latter manner that the fine chrome green used for painting on porcelain is ob- 
tained. Chromium sesquioxide is not decomposed by heat, and fuses only at 
very elevated temperatures. A corresponding hydroxide, Cr(OH) 3 , is thrown 
down as a bluish-green precipitate when ammonia- water is added to the green 
solution of chromic chloride. The same hydroxide is precipitated by the alka- 
line hydroxides, but dissolves in an excess of the reagent ; when the liquid 
is boiled, an insoluble hydroxide is thrown down. 

546. Chromic Anhydride. CrO 3 . — This compound is com- 
monly called chromic acid. It is prepared by mixing a cold satu- 



326 LESSONS IN CHEMISTRY. 

rated solution of potassium dichromate with one and a half times 
its volume of strong sulphuric acid. As the liquid cools, chro- 
mium anhydride separates in crimson needles, which are quickly 
drained on a dry brick and recrystallized in the smallest possible 
quantity of warm water. It is a deliquescent substance, exceed- 
ingly soluble in water, and the solution has an orange color. It 
energetically oxidizes many bodies. With hydrochloric acid it 
forms water and chromic chloride, while chlorine is set free. 
2Cr0 3 + 12HC1 = 2CrCl 3 + 6H20 + 3C1 2 

It instantly oxidizes sulphur dioxide, chromium sulphate being 

formed. 

2Cr0 3 + 3S0 2 = Cr 2 (SO±) 3 

It oxidizes alcohol and ether with such energy that those com- 
pounds are inflamed. 

547. Chromates. — The solution of chromium anhydride must 
be regarded as containing chromic acid, H 2 Cr0 4 = H 2 + CrO 3 , 
corresponding in molecular constitution to sulphuric acid. The 
chromium compounds are all derived from potassium dichromate, 
which is manufactured from chrome iron. A mixture of the 
pulverized mineral with potash and lime is roasted with full 
access of air : calcium and potassium chromates are produced, 
oxygen being absorbed. The mass is extracted with a hot solution 
of potassium sulphate, which converts the calcium chromate 
into the potassium salt. The decanted yellow liquid is then 
mixed with enough sulphuric acid to form the dichromate. 

2K 2 CrO* + H2S0* = K 2 Cr 2 07 + K 2 SO* + IPO 
This compound, being less soluble than the neutral chromate, 
crystallizes as the solution cools. 

548. Potassium chromate, K 2 Cr0 4 , forms beautiful, lemon-yel- 
low, anhydrous crystals, which are very soluble in water, to which 
they impart an intense yellow color. 

549. Potassium dichromate, K 2 Cr 2 T , forms large, orange-red 
crystals, soluble in about eight times their weight of cold water, 
and in much less boiling water. By heat they are decomposed 
into potassium chromate, chromium oxide, and oxygen. 

2K 2 Cr 2 7 = 2K 2 Cr0 4 + Cr 2 3 + O 3 



CHROMIUM. 327 

Potassium dichromate is an energetic oxidizing agent. When 
heated with sulphuric acid, it yields oxygen, while chrome alum 
will crystallize from the liquid obtained by treating the residue 
with boiling water. 

K 2 Cr 2 7 + 4H 2 SO* = Cr 2 (SO±) 3 .K 2 S0 4 -f 4H 2 + O 3 

When sulphur dioxide is passed into a solution of potassium 
dichromate, the orange color is gradually replaced by green : while 
the sulphur dioxide becomes sulphuric acid, both chromium and 
potassium are converted into sulphates, and the liquid will yield 
chrome alum if sulphuric acid be added. 

K 2 Cr 2 7 + 3S0 2 + H 2 SO± = Cr 2 (S0 4 ) 3 .K 2 SO* + H 2 

550. Ammonium dichromate, (NH 4 ) 2 Cr'0 7 , may be made by dividing a solu- 
tion of chromic acid into two equal portions, neutralizing one with ammonia- 
water, and then adding the other. When the solution is evaporated, the am- 
monium dichromate separates in red crystals, which, when heated, yield pure 
chromium trioxide in a curious pulverulent form, resembling green tea. 

(NH*) 2 Cr 2 7 = Cr 2 3 + 4H 2 + X 2 

551. Lead chro/nate, PbCrO 4 , occurs in the mineral crocoite, and is made 
by mixing solutions of potassium chromate and lead acetate; potassium ace- 
tate remains in solution, while the lead chromate forms a dense yellow pre- 
cipitate, which when washed and dried constitutes chrome yellow. It is 
insoluble in water and in acetic acid, but dissolves in solutions of the alkaline 
hydroxides. It melts at a red heat, and is readily reduced by both hydrogen 
and charcoal. It is sometimes substituted for cupric oxide in the analysis of 
carbon compounds. 

552. Tests for Chromium. — Although there is a series of 
chromous salts in which the chromium atom is diatomic, the 
reactions of the salts corresponding to the sesquioxide are suffi- 
cient to characterize this element. In the green solutions of these 
compounds hydrogen sulphide produces no precipitate : ammonium 
sulphide throws down chromium hydroxide, hydrogen sulphide 
being disengaged. 

2CrCl 3 + 3(NH 4 ) 2 3 + 6H 2 = 2Cr(OH) s + 6XH 4 C1 + 3H 2 8 
The alkaline hydroxides and ammonia produce the same green pre- 
cipitate, which dissolves readily in an excess of the former re- 
agents, more slowly in ammonia-water. When the solution so 
obtained is boiled, anhydrous chromium sesquioxide is precipitated, 
and does not redissolve on cooling. 



328 



LESSONS IN CHEMISTRY. 



The chromates may be identified by heating them in a test- 
tube with a little common salt and sulphuric acid. Irritating 
red vapors of chromyl chloride, Cr0 2 G 2 , are disengaged, and if 
conducted into a cold tube will condense to a blood-red liquid. 
When passed into water, these vapors are decomposed into 
hydrochloric and chromic acids. 

Cr0 2 C12 + IPO = Cr03 + 2HC1 
In neutral solutions of chromates, barium chloride and lead 
acetate produce yellow precipitates, and silver nitrate throws 
down a red compound. 

553. Closely related to chromium in their chemical relations are the elements 
molybdenum, tungsten, and uranium. Molybdenum occurs chiefly as the sul- 
phide, MoS 2 , in the mineral molybdenite. Tungsten is found in various tung- 
states, as in the mineral wolfram, which is a tungstate of iron and manganese. 
The principal uranium mineral is pitchblende, an impure oxide of uranium of 
varying color. It was in a variety of this mineral that helium was discovered. 
The three metals have been recently produced in considerable quantities and 
in a very pure state by means of the electric furnace. They resemble chro- 
mium in their properties : they form trioxides, which, like chromium triox- 

ide, are the anhyrides of corresponding 
acids. A number of lower oxides are 
also known. 

554. Tin, Sn = 118.— Tin is 
rarely found in the metallic state 
in nature : its only workable ore 
is the dioxide, which constitutes 
the mineral cassiterite. The 
principal tin-mines are in Corn- 
wall, England, and in various 
parts of Farther India and 
Australia. 

The ore is crushed, and the dioxide, 
being very heavy, can be separated 
PlG. 1 24. from the lighter earthy matters by wash- 

ing in a stream of water. It is then 
roasted, the sulphides and arsenides of iron and copper present being converted 
into oxides, which are removed by a second washing. The purified cassiterite 
is mixed with charcoal and fed into a cupola furnace (Fig. 124), where the 
combustion is supported by a blast of air. The reduced tin collects on the 
hearth of the furnace, and runs into a basin, where it is stirred with poles of 




tin. 329 

green wood. The gases given off reduce any oxide that has been formed, and 
bring to the surface of the molten metal the foreign matters, which form a 
dross. The tin is further purified by being melted at a low temperature on the 
inclined hearth of a reverberatory furnace. Being more fusible than the foreign 
metals present, it runs into a cavity prepared for it, while the less fusible 
metals remain on the hearth. 

Tin is a silvery-white metal, having a density of about 7.3. It 
melts at 228°, and may be crystallized by slow cooling. It is 
malleable and ductile : when a bar of tin is bent, it produces a 
peculiar noise, called the cry of tin, caused by the sliding of the 
crystals over one another. 

It is not affected by the air at ordinary temperatures, but when 
melted absorbs oxygen, and by stirring may be entirely converted 
into the dioxide. It is dissolved by hydrochloric acid, hydrogen 
being disengaged, while stannous chloride is formed. Nitric acid 
converts tin into dioxide, giving off torrents of red vapors. Hot 
solutions of the alkaline hydroxides dissolve tin, forming alkaline 
stannates, and disengaging hydrogen. 

Tin is used for the manufacture of tin foil, employed for 
enveloping tobacco, chocolate, etc. ; also for tinning copper and 
iron, which is accomplished by dipping the perfectly clean objects 
into a bath of molten tin. Its resistance to the action of vegetable 
acids renders it invaluable as a coating for culinary utensils. It 
enters into the composition of plumbers' solder, which is an alloy 
of tin and lead. Bronze, bell-metal, gun-metal, and speculum- 
metal are alloys of tin and copper (page 287). Britannia metal 
is tin alloyed with a small proportion of antimony, bismuth, and 
copper. 

In its compounds tin is either diatomic or tetratomic. Those 
in which it is diatomic are called stannous compounds, while in 
the stannic compounds it is tetratomic, one atom of tin having the 
same combining power as four atoms of hydrogen. 

555. Stannous Chloride, SnCl 2 . — Anhydrous stannous chlo- 
ride is obtained by passing hydrochloric acid gas over heated tin. 
It is a white solid, fusible at 250°. When it is dissolved in a 
small quantity of water, or when metallic tin is dissolved in hot 
hydrochloric acid, a solution is obtained which when sufficiently 
concentrated deposits crystals of a hydrate containing SnCl 2 + 



330 LESSONS IN CHEMISTRY. 

2H 2 0. These crystals are known in commerce as tin salt or tin 
crystals. They are soluble in a small quantity of water, but when 
the solution is diluted, a deposit of an oxychloride is formed con- 
taining SnCP.SnO. At the same time a certain proportion of 
the stannous chloride is converted into stannic chloride, SnCl 4 . 
This decomposition is prevented by the presence of free hydro- 
chloric acid, or by a small quantity of ammonium chloride. Stan- 
nous chloride is a reducing agent: a few drops of its solution 
instantly decolorize the purple solution of potassium permanga- 
nate ; it reduces the salts of silver and gold, setting free the metal. 
When it is added to a solution of mercuric chloride, a white pre- 
cipitate of mercurous chloride is formed, which an excess of 
stannous chloride converts into a gray deposit of finely-divided 
metallic mercury. In these reactions the stannous chloride be- 
comes stannic chloride. Stannous chloride is used as a mordant 
in dyeing. 

556. Stannic Chloride, SnCl 4 , is formed with the production of light and 
heat by the direct union of tin and chlorine. It is prepared by passing dry 
chlorine over melted tin contained in a retort ; it then distils, and condenses 
as a heavy, fuming, yellow liquid, which boils at 120°. It combines energeti- 
cally with water, forming crystals of a hydrate containing SnCl 4 -f 5 IPO. The 
same hydrate may be made by dissolving tin in hydrochloric acid and from 
time to time adding small quantities of nitric acid. The crystals are soluble 
in water, yielding a limpid solution. 

557. Stannic Oxide, SnO 2 . — When an alkaline hydroxide is added to a solu- 
tion of stannous chloride, stannous hydrate is formed as a white precipitate, 
which, by boiling, is converted into black stannous oxide, SnO. The addition 
of ammonia to a solution of stannic chloride throws down a white gelati- 
nous precipitate of stannic hydrate, H 2 Sn0 3 , which by the action of heat is 
converted into stannic oxide, SnO 2 . This compound is found in nature in 
hard, transparent crystals ; it is cassiterite. It is an acid oxide, and stannic 
hydrate reacts with the bases, forming stannates whose compositions corre- 
spond to the carbonates. The white powder produced by the action of nitric 
acid on tin is a stannic hydrate, containing Sn(OH) 4 = SnO 2 + 2H 2 0. 

558. Sulphides of Tin. — By heating together the proper proportions of tin 
and sulphur, two sulphides may be obtained. Stannous sulphide, SnS, is a 
gray, crystalline mass. The preparation of stannic sulphide, SnS 2 , requires 
particular precautions ; an amalgam of tin with half its weight of mercury 
is mixed with flowers of sulphur and ammonium chloride and heated to dull 
redness. Mercuric sulphide, ammonium chloride, and the excess of sulphur 
sublime, while the interior of the vessel becomes lined with a golden-yellow, 



PLATINUM AND ITS ALLIED METALS. 331 

crystalline mass of stannic sulphide. The operation must then be arrested, 
or this compound will be decomposed into stannous sulphide and sulphur; the 
addition of the mercury and ammonium chloride is intended to keep the tem- 
perature down to the volatilizing points of mercuric sulphide and ammonium 
chloride. Stannic sulphide forms soft, crystalline scales, called mosaic gold. 

559. Tests for Tin. — In stannous solutions, both hydrogen 
sulphide and ammonium sulphide form brown precipitates, soluble 
in yellow ammonium sulphide (§ 143). The alkaline hydroxides 
and ammonia give white precipitates, soluble in an excess of the 
former reagents, but insoluble in ammonia. Gold trichloride throws 
down purple of Cassius. In mercuric chloride solutions, an excess 
of stannous chloride precipitates gray metallic mercury. 

In stannic solutions, hydrogen sulphide and ammonium sulphide 
form yellow precipitates, soluble in a large quantity of the latter 
reagent. Apiece of zinc placed in either a stannous or a stannic 
solution becomes covered with a deposit of tin, which may be 
rendered brilliant by burnishing. 

560. The elements titanium, zirconium, and thorium closely resemble tin in 
their chemical relations, although their physical properties are very different. 
Each forms a tetrachloride and a dioxide. Titanium occurs as dioxide in the 
minerals rutile, anatase, and brookite. Zirconium exists as a silicate in zircon, 
and thorium is a constituent of a number of minerals, such as ruonazite and 
thorite. The oxides of thorium and zirconium have found an application in 
the " Welsbach" burner, in which a gauze of these compounds is rendered 
incandescent by a non-luminous gas flame. 

The very rare element germanium also belongs to this group. 



LESSON LXII. 

PLATINUM AND ITS ALLIED METALS. 

561. Like gold, platinum is found in the metallic state in 
rounded granules distributed through sandy deposits. Being 
very heavy, it is also separated like gold, by washing the sand in 
a stream of water. The native platinum, however, is not pure : 
besides containing traces of gold, copper, and iron, it is alloyed 
with several other metals which it resembles in certain properties, 
and which are called the platinum metals. They are rhodium, 



332 LESSONS IN CHEMISTRY. 

ruthenium, palladium, iridium, and osmium. The platinum is ex- 
tracted by treating the grains first with dilute nitre-hydrochloric 
acid, which removes all excepting the platinum metals, and then 
heating it with strong nitro-hydrochloric acid, which dissolves the 
platinum, leaving osmium and the greater part of the iridium. 
The liquid is then exactly neutralized with sodium carbonate, and 
a solution of mercuric cyanide is added. This throws down a 
precipitate of palladium cyanide, which is removed by filtration, 
and the clear liquid is treated with ammonium chloride. A crys- 
talline precipitate of a double chloride of platinum and am- 
monium forms, and this, when calcined, leaves a porous gray 
residue of platinum sponge. Platinum so prepared always con- 
tains some iridium, for the latter metal also separates as a double 
chloride when the ammonium chloride is added. 

In order to agglomerate the spongy platinum, it is made into 
a stiff paste with a little water, and this is strongly compressed 
in a slightly conical steel cylinder. It is then removed, heated to 
whiteness, and converted into a solid mass by hammering. Plati- 
num is also melted in lime crucibles, heated by the flame of the 
oxyhydrogen blow-pipe directed against the mass of metal. 

Platinum is a grayish-white, lustrous metal. Its density is 
21.5. It is very malleable and ductile. It melts in the oxy- 
hydrogen flame and in the electric furnace ; at a white heat it 
becomes soft and can be forged and welded like iron. It is not 
affected by the air at any temperature, and does not dissolve in 
either hydrochloric, sulphuric, or nitric acid. When alloyed 
with silver, it is attacked by nitric acid. Nitro-hydrochloric 
acid dissolves it slowly in the cold, more rapidly by the aid 
of heat, converting it into the tetrachloride. The alkaline hy- 
droxides and nitrates attack it at a high temperature, and these 
substances must not be fused in platinum crucibles. 

Platinum has the power of condensing gases in its pores, and 
we have already seen how the oxidation of ammonia and vapor 
of alcohol and ether may be effected by a platinum wire. This 
property is more strongly manifested by platinum sponge than 
by the compact metal, and hydrogen escaping from a jet may 
be ignited by holding in it a morsel of recently-heated spongy 



PLATINUM AND ITS ALLIED METALS. 333 

platinum. When a solution of platinic chloride is boiled with 
potassium hydroxide, and alcohol is added to the boiling liquid 
with constant stirring, metallic platinum is deposited as a black 
powder. This powder, which is called platinum-black, is in a 
state of extreme division, and brings about the oxidation of com- 
bustible gases and vapors even more readily than platinum sponge. 

Platinum is employed for the manufacture of crucibles and 
dishes for the laboratory, for it not only resists high temperatures 
but is attacked by very few chemical reagents. It is manufac- 
tured into large retorts for the concentration of sulphuric acid. 
All of this apparatus is made as thin as is consistent with strength, 
for the metal is costly. 

Platinum forms two series of compounds, — platinous compounds, 
in which it is diatomic, and platinic compounds, in which it is 
tetratomic. 

562. Platinic Chloride, PtCl 4 , is made by dissolving the 
metal in nitro-hydrochloric acid. When the reddish-brown liquid 
is sufficiently concentrated, it deposits, on cooling, hydrated crys- 
tals of platinic chloride, which may be rendered anhydrous by 
heat. The anhydrous salt is a red-brown deliquescent mass, very 
soluble in water, alcohol, and ether. With hydrochloric acid 
it forms the compound PtCl 4 .2HCl, which crystallizes with six 
molecules of water, and is called cliloro -platinic acid. The aque- 
ous solution of this compound produces yellow crystalline precipi- 
tates in solutions of potassium and ammonium chlorides. They 
are chloroplatinates, and have the compositions PtCl 4 .2KCl and 
PtCl 4 .2NH 4 Cl. They are slightly soluble in cold water, and in- 
soluble in alcohol, but dissolve readily in boiling water. When the 
ammonium salt is heated, it leaves a residue of spongy platinum. 

If platinum tetrachloride be carefully heated to 200°, chlorine 

is disengaged; and if the residue be extracted with boiling water, 

the unaltered platinic chloride is dissolved, while platinum dichlo- 

ride, PtCP, remains as an olive-green powder. 

Of the other metals of the platinum group, Osmium is the least fusible : it 
has never been melted. "When strongly heated in the air, it forms a volatile 
oxide, which is one of the most dangerous poisons known. It is the heaviest 
known element, having a density of 22.48. 

563. Palladium is the most fusible of these metals, and has the lowest 



334 LESSONS IN CHEMISTRY. 

density, its specific weight being 11.4. When a piece of this metal is made 
the negative electrode of an apparatus in which water is being decomposed 
by the voltaic current, it will absorb about nine hundred times its volume of 
hydrogen. 

564. Iridium often constitutes a considerable proportion of platinum ore, 
which is then called platiniridium or osmiridium, as platinum or osmium pre- 
ponderates in the alloy. Iridium alloyed with 90 per cent, of platinum is as 
hard and elastic as steel, is less fusible than platinum, and is unaltered by the 
air. It is used for the points of gold pens and ink-containing pencils, as is 
also the native alloy, osmiridium. The density of iridium is 22.38. 

565. Rhodium is more fusible than iridium, but less fusible than platinum. 
Its density is 12.1. It does not dissolve in nitro-hydrochloric acid unless it is 
alloyed with other metals. 

566. Ruthenium is the most infusible metal after osmium. Its density is 
12.26. It is hardly attacked by boiling nitro-hydrochloric acid. 



LESSON LXIII. 

THE CHEMISTRY OF LIFE. 

567. Under the influence of the mysterious principle which 
we call life, certain chemical compounds undergo a complete meta- 
morphosis. Their elements become rearranged in manners which 
are beyond our methods of research, and the matter becomes or- 
ganized. It assumes certain definite forms which we call cells, 
and living cells are gifted with a wonderful power of reproduc- 
tion : under the proper conditions they can convert unorganized 
dead matter into other cells, either of the same kind or of very 
different kinds related by progressive modifications. Chemists 
are able to change one form of matter into another, — to modify 
and destroy molecules, and to construct new molecules ; they are 
unable to create the simplest cell, the lowest form of organized 
matter. We have no reason to believe tbat any cell is ever pro- 
duced except from another, but we know that under modified cir- 
cumstances the nature of the cell may in the course of succes- 
sive reproductions become completely modified, and new forms of 
organized matter are produced. 

The elements which enter into the composition of organized 



THE CHEMISTRY OF LIFE. 335 

matter are comparatively few: all vital tissues contain carbon, 
hydrogen, and oxygen, and these three, together with nitrogen 
and a few salts, principally phosphates, chlorides, and sulphates of 
sodium, potassium, and calcium, constitute the greater part of all 
tissues, vegetable and animal. Plants directly convert carbon 
dioxide and water into carbohydrates, such as cellulose, starch 
and the sugars, and other compounds, even including hydrocar- 
bons. In this reducing action of vegetable life on carbon dioxide 
and water, the atoms of carbon and hydrogen recover part of the 
energy which disappears from them in the formation of those 
compounds. As far as their matter is concerned, plants then act 
as storers or regenerators of energy. In the natural heat and 
motion of animals the atomic energy of the compounds of carbon 
and hydrogen is manifested as those compounds are again oxidized 
with the formation of carbon dioxide and water. Animal life is 
really dependent on the continual expenditure of energy. Vege- 
tables can receive their nutrition directly from mineral matter, 
but animals can form tissues only from matter that has first been 
prepared by vegetables or other animals. 

However, besides the carbon, hydrogen, and oxygen, plants 
absorb nitrogen from nitrogenized matters in the soil, and the 
nitrogen compounds of plants are essential for the nutrition of 
animals. We must study some of the compounds which have 
thus far been formed only under the influence of life, and of 
which the organization results in the production of cells. Among 
these substances, which we must remember are not chemical com- 
pounds of definite and known constitution, of first importance are 
the albuminoid matters. 

ALBUMINOID AND GELATINOID SUBSTANCES. 

568. These complex matters are composed of carbon, hydro- 
gen, oxygen, and nitrogen, and a small proportion of sulphur. 
By their compositions and properties, they are all related to the 
albumen of white of egg, or to the gelatin or glue which can be 
extracted from bones. If flour made from wheat or other cereal 
be kneaded in water, the starch is washed out, while a gray elastic 



336 LESSONS IN CHEMISTRY. 

mass of gluten remains. This gluten may be separated into sev- 
eral different substances, having different degrees of solubility in 
alcohol : they are of similar composition, containing a little more 
than 50 per cent, of carbon, 7 of hydrogen, 17 of nitrogen, 20 of 
oxygen, and rather less than one per cent, of sulphur. The water 
used in the preparation of gluten contains another matter, which 
may be separated by allowing the starch to settle, adding a few 
drops of acid to the clear liquid, and heating to the boiling point. 
An albuminoid matter then coagulates in white flakes. Its com- 
position does not differ greatly from that of the bodies separated 
from gluten, but there are slight differences depending on the 
grain or seed from which the substance is derived. From the 
seeds of leguminous vegetables, such as peas, beans, and lentils, a 
body called legumine may be extracted, and, in addition to the ele- 
ments contained in gluten and other vegetable albumens, this sub- 
stance contains a small percentage of phosphoric acid, probably in 
the form of a substituted acid, in which various carbon groups 
replace one or more hydroxyl groups of orthophosphoric acid. 

The albuminoid matters of animals, which are derived from the 
similar vegetable substances, are classified more with reference to 
their behavior under the action of heat, and their solubilities in 
water, acids, alkalies, etc., than according to their composition, 
which varies but little. They may, however, be arranged in 
two groups, — albuminoid matters and gelatin-like compounds. 

The general composition of these bodies is as follows : 

Albumen Group. Gelatin Group. 

Carbon 53.5 50.0 

Hydrogen 6.9 6.6 

Oxygen . 23.0 26.1 to 23.1 

Nitrogen 15.6 16.8 

Sulphur 1.0 0.5 to 3.5 

Of the albuminoid matters we can consider only albumin, 
fibrin, casein, and hemoglobin. 

569. Albumin exists in a soluble form and an insoluble modi- 
fication. Soluble, it occurs in white of egg, and in the serum or 
clear liquid of blood ; but even these forms present certain differ- 
ences. If either of these liquids be evaporated at a low temper- 



ALBUMINOID BODIES. 337 

ature, the albumin remains as a transparent, yellowish, gum-like 
mass, which is perfectly soluble in water. It is not pure, but 
contains a small quantity of alkaline carbonate and certain salts. 
If a solution of albumin be heated to 70°, it becomes clouded, 
and at a few degrees higher the albumin separates either in flakes 
or in a solid mass, according to the concentration of the solution. 
The soluble albumin has coagulated and has become insoluble 
albumin. Solutions of albumin are also coagulated by the addi- 
tion of either sulphuric, nitric, or hydrochloric acid, or of certain 
salts, such as mercuric chloride and lead acetate. Metaphosphoric 
acid instantly precipitates albumin from its solutions. Ortho- 
phosphoric acid, acetic and lactic acids, form no precipitates with 
albumin, neither does common salt unless acetic acid be present. 

570. Fibrin. — When fresh blood is allowed to stand, it soon 
separates into a yellow liquid, called serum, and a red coagulum or 
clot. The clot contains the red corpuscles, the oxygen -carriers of 
the blood, imprisoned in a mass of insoluble albuminoid matter. 
By beating the fresh blood with a bunch of twigs or an egg- 
beater, the mass of blood is prevented from coagulating, and the 
albuminoid matter, which is called fibrin, becomes attached to the 
beater in red flakes. By washing in a stream of water, the red 
corpuscles are washed out, and the fibrin remains as light-gray, 
elastic filaments. It is insoluble in water, but dissolves in very 
dilute alkaline solutions. Fibrin is formed by the union of two 
substances contained in the blood, whenever that liquid is kept at 
rest. Its spontaneous coagulation causes the cessation of bleeding 
from slight cuts and other small wounds. 

The stiffening of the muscles which takes place soon after death, 
is due to the coagulation of a peculiar albuminoid matter, called 
myosin, which exists in solution in the muscular tissues. It is 
soluble in water containing 10 per cent, of salt, but is precipitated 
by a larger quantity : it is extracted by virtue of this property. 

571. Hemoglobin is a crystallizable matter which can be 
extracted from the red corpuscles of blood. It contains a small 
proportion of iron. Hemoglobin has the property of absorbing 
oxygen and forming an unstable compound from which the oxygen 

22 



338 LESSONS IN CHEMISTRY. 

escapes by exposure in a vacuum. It is probably by this property 
of the hemoglobin which they contain, that the red blood cor- 
puscles are enabled to carry oxygen to all parts of the system. 
Hemoglobin will also absorb carbon monoxide, and when it has 
absorbed that gas it is incapable of combining with oxygen : this 
explains the poisonous effects of carbon monoxide on the system. 
Hydrogen sulphide reduces hemoglobin, — that is, removes its 
oxygen ; and we can so understand the injurious action of any 
quantities of this gas. 

572. Casein, Milk. — Milk is a dilute solution of lactose or 
milk sugar and a small quantity of mineral salts, in which are 
suspended very small fat globules, and, either suspended or in 
solution, an albuminoid matter called casein. The specific gravity 
of milk is about 1.030. After standing for several hours, the 
greater number of the fat globules come to the surface, consti- 
tuting the cream ; cream, however, contains some lactose, casein, 
and salts. Its composition varies greatly, as may be seen from the 
following results of the analysis of three samples : 

i. ii. in. 

Water 72.2 66.36 50 

Fat 19.0 18.87 43.9 

Casein, lactose, salts 8.8 14.77 6.1 

The following is the composition of an average sample of cow's 
milk, but it must be borne in mind that no two samples will prob- 
ably have exactly the same composition : 

Water 87.0 per cent. 

Fat 3.5 

Lactose 4.8 " 

Casein 4.0 " 

Salts 0.7 " 

When an acid is added to milk, a thick deposit of coagulated 
casein is formed. This same precipitation occurs when milk 
naturally becomes sour by the formation of lactic acid. Casein 
closely resembles insoluble albumin ; it is insoluble in water, 
but dissolves in dilute solutions of the alkaline hydroxides and 
carbonates, and it is probably in solution in fresh milk, for that 
liquid has an alkaline reaction. Casein is the characteristic 
constituent of cheese. 



THE CHEMISTRY OF LIFE. 339 

In cheese-making, rennet, a substance obtained from the fourth 
stomach of the calf, is added to the milk : this precipitates the 
casein, and the latter, enveloping the fat-globules, carries them 
down with it. The curd, separated from the liquid, constitutes 
cheese. 

573. Gelatin. — When bones are immersed in hydrochloric 
acid, the mineral matter, consisting principally of calcium phos- 
phate and carbonate, is dissolved, and a semi-transparent, elastic 
mass is obtained, retaining the form of the bone. This body is 
insoluble in cold water, but by long boiling it dissolves, and, on 
cooling, the solution sets in a transparent jelly. This substance 
is gelatin, or glue. It is not peculiar to the bones, but exists also 
in certain other tissues, particularly in the skin, and in the swim- 
ming-bladders of fishes. The best gelatin is obtained from the 
swimming-bladder of the sturgeon ; it is called fish-glue. Very 
little is known regarding the difference between gelatin and the 
substances from which it is derived. 

Dry gelatin occurs in transparent or translucent sonorous sheets, 
whose color varies from colorless to brown, according to the purity. 
It swells in water, but does not dissolve until the liquid is heated. 
Its solution is precipitated by alcohol, but not by acids, with the 
exception of tannic acid, with which it forms an insoluble com- 
pound. The tanning of skins and hides depends on the formation 
of this compound in the body of the skin, which is so converted 
into leather. 

574. By the processes of digestion, the vegetable and animal 
matters which serve as food are converted into substances which 
can be assimilated or made part of our bodies. These processes 
begin in the mouth, where the starchy substances encounter in 
the saliva a peculiar unorganized ferment called ptyalin, which is 
capable of transforming them into soluble glucose. Ptyalin is 
probably identical with diastase, which is formed during the germi- 
nation of grain. In the stomach, the conversion of starch into 
glucose continues, and the albuminoid matters are converted into 
soluble bodies called peptones by another ferment, pepsin, contained, 



340 LESSONS IN CHEMISTRY. 

together with a little hydrochloric acid, in the gastric juice. This 
ferment exists in rennet, obtained from the stomach of the calf, 
and used in the manufacture of cheese. The peptones appear to 
be formed by the hydration of the albuminoid bodies, and in 
the system they are probably converted into all the varieties of 
albuminoid tissue. As the food passes from the stomach it 
encounters in the small intestines other ferments, by which the 
fatty matters are emulsified and rendered capable of being ab- 
sorbed and passing into the blood, by which they are carried 
and deposited where needed in the system. 

The slow combustion by which life is sustained results in the 
oxidation of the tissues, and the removal of the matters no longer 
useful. This oxidation is not accomplished in one operation, but 
in several stages, during which many compounds intermediate 
between the albuminoid and fatty bodies, and the carbon dioxide, 
water, and nitrogen which would result from their complete com- 
bustion, are formed. We have seen that a great part of the carbon 
and hydrogen is indeed removed as carbon dioxide and water, but 
the salts are in great part eliminated unchanged by the urine and 
the perspiration. The nitrogen is excreted principally as urea, 
phosphorus as sodium acid phosphate, sulphur as sodium sulphate, 
etc. A small part of the nitrogen of the system is excreted in 
forms intermediate between the albuminoid bodies and urea. 
Among the more important of these is uric acid, C 5 H 4 N 4 3 , a 
compound forming a small proportion of human urine, and exist- 
ing in large quantity in the solid urine of birds and reptiles. 

In all these processes there is comparatively little that we can 
understand. We know only that they all result in a transfer of 
energy, — that in living matter the chemical energy is converted 
into the energy of Jife ; and we can comprehend only the beginning 
and the end of the phenomenon, — the forms of matter which are 
capable of organization, and the products of the disorganization 
without which life could not continue. 



APPENDIX. 



i. 

CKYSTALLOGKAPHY. 

A crystal is a natural polyhedron ; that is, a solid bounded by plane 
surfaces or faces. The greater number of solid chemical substances 
form more or less perfect crystals whenever a certain freedom of mo- 
tion is communicated to their molecules, so that these molecules may 
arrange themselves without interference. Such freedom of motion 
may be given to the molecules : 

1. By dissolving the solid in any liquid by which it is not altered, 
and allowing a hot saturated solution to cool slowly, or by the spon- 
taneous evaporation of the solvent if the solid be equally soluble at 
all temperatures. Potassium chlorate, lead iodide, potassium nitrate, 
and alum may be crystallized from water by the first method ; com- 
mon salt by the second. 

2. By melting the solid, and decanting the still liquid portion after 
a crust has formed on its surface. Sulphur and bismuth may be so 
crystallized. 

3. By subliming the solid, and allowing the vapor to condense 
very slowly. In this manner fine crystals of iodine and camphor 
may be obtained. 

4. By a chemical reaction in which the crystallizable substance is 
formed in a medium in which it is insoluble. Potassium chloro- 
platinate and potassium acid tartrate are so formed in microscopic 
crystals. 

A solid which manifests no tendency to become crystalline is said 
to be amorphous. Glass and glue are examples of amorphous sub- 
stances. 

For the sake of convenience crystals are classified in six si/stems, and 

341 



342 



APPENDIX. 



each system is characterized by a set of axes, which are imaginary 
straight lines passing through the centre of the crystal and joining 
opposite solid angles or the centres of opposite faces or edges. The 
forms belonging to any one system may be derived from each other 
by replacing the edges or angles by plane surfaces ; in all such deriva- 
tives the imaginary axes must remain unchanged. 

1. The Isometric System has three equal axes at right angles to 
each other. The type of the system is the octahedron (a), in which 
the axes join the opposite angles. The tetrahedron (b) is derived from 





this form by extending the alternate faces until they meet, forming 
edges of which the centres are then joined by the axes. The cube (c) 
is obtained by replacing or truncating the six angles of the octahedron 
by as many faces, and extending these until they form edges; the 
axes join the centres of opposite faces. By replacing or bevelling the 



I 



00 




twelve edges of the octahedron by the same number of faces so that 
these meet in the centres of the octahedral faces, the rhombic dodeca- 
hedron (d ) is produced. 

Besides these simpler forms, the isometric system includes a number 
of others which are bounded by twelve, twenty-four, and even forty- 
eight faces. 



APPENDIX. 



343 



2. The Tetragonal System is characterized by three axes inter- 
secting at right angles, of which two are equal and the third either 





longer or shorter. The type is the tetragonal pyramid (e). in which 
the axes join the opposite angles. A similar form (/), in which two 
of the axes connect the centres of opposite edges, while the third axis 



--K-n- 



1— 

1 


1 

N _ 
"" !\ 1 

1 


I ' \ 



(9) 



(A) 



joins the vertices, is called a pyramid of the second order. Prisms 
of the first (g) and of the second (/*) order, as well as more complex 
forms, occur in this system. 

3. The Orthorhombic System has three unequal axes at right 
angles to each other. The typical form is the orthorhombic pyramid 
(i), in which the axes connect the opposite angles. Other forms of 
this system are designated as prisms (&), domes, and pinacoids. 



344 



APPENDIX. 



4. The Hexagonal System has four axes ; three are in the same 
plane, equal in length, and intersect at angles of 60° ; the fourth is at 




(*) 



right angles to these three, and may be either longer or shorter. In 
the typical form, the hexagonal pyramid of the first order (£), the four 





axes join the solid angles ; in the pyramid of the second order (ra), the 
three equal axes connect the centres of opposite edges, while the fourth 






; 


^L> 




■ i 
! i 

1 -4c-- 




k 


■- \ 1 - 

1 1 
1 1 

"~ 1 


— k >. 



(°) 



axis joins the two vertices. The commonest form of this system, the 
rhombohedron {n) ) results from the pyramids by extending their 



APPENDIX. 



345 



alternate faces till they meet : it is bounded by six rhombus-shaped 
faces. 

Prisms of the first (o) and second order (jo), and other more com- 
plex forms, belong to this system. 

5. The Moxoclinic System has three unequal axes, of which two 
are obliquely inclined to each other in a plane intersecting the third 
axis at right angles. The plane containing the first two axes divides 
the forms of this system into symmetrical halves : monoclinic crystals 
are characterized by their bilateral symmetry. 

Monoclinic pyramids (q) have two sets of faces. There are also 
prisms, domes, and pinacoids. 



< 

< 


J f 

f 


1 > 


/ 



(p) 




6. The Trtclinic System has three unequal axes which are ob- 
liquely inclined to each other. 

The names of the triclinic forms are analogous to those of the or- 
thorhombic and monoclinic systems. 

The crystal forms enumerated above are either closed forms or open 
forms. The former may exist by themselves, as, for example, a cube 
or a rhombohedron, while the latter require one or several other forms 
to enclose space. Most of the natural crystals show combinations of 
several forms, modifying each other in various manners. The char- 
acteristic axes, however, are not affected by these modifications. 

A substance which crystallizes in forms belonging to two different 
systems is said to be dimorphous. Such a body is sulphur. 

Two different substances whose crystals are of precisely the same 
form are said to be isomorphous. The chlorides, bromides, and 
iodides of potassium and sodium are isomorphous ; the alums and the 
spinels are also excellent examples of isomorphism. 



346 APPENDIX. 

II. 

STEKEOCHEMISTKY. 

Of the very numerous cases of isomerism which have been observed 
among the carbon compounds, the great majority may be explained by 
assigning different molecular structure to the different compounds of 
the same composition. Structural formulae are employed to represent 
the relations supposed to exist between the atoms constituting the 
molecules. A few instances, however, have long been known of com- 
pounds whose constitution must be expressed by the same structural 
formula, although they differ in certain of their properties. The 
lactic acids (see p. 216) afford an example of this mode of isomerism. 
Ordinary lactic acid, which is produced in the lactic fermentation of 
sugar, when dissolved in water has no action upon polarized light, 
while solutions of the other two modifications have the property of 
rotating the plane of polarization : sarcolactic or dextrolactic acid de- 
flects it to the right, and levolactic acid to the left. In the formula 
CH 3 .CH(OH).COOH, one of the carbon atoms is combined with an 
atom of hydrogen, and with one of each of the radicals hydroxyl, 
methyl, and carboxyl ; the four atomicities of this carbon atom are em- 
ployed in uniting four different atoms or groups with it. 

H 

I 
CH 3 — C— OH 

COOH 

It has been shown that all the other known carbon compounds which 

are optically active (rotate the plane of polarized light) contain one 

or more such carbon atoms, and further, that these compounds always 

* exist in several modifications. To account for these remarkable facts 

the following theory has been proposed. 

All the known facts indicate that the four atomicities of the carbon 
atom are exactly alike. Now, if we conceive an atom to be so situated 
as to form the centre of an isometric tetrahedron, and the four atoms or 
groups with which the atom is united, at the vertices of this tetrahe- 
dron, the affinities would be symmetrically distributed. 

As long as any two atoms or groups of the same kind occupy posi- 
tions at the vertices, only one form is possible ; * but when the four 



* This can only be shown by means of stereochemical models. 



APPENDIX. 



347 



atoms or groups are all different, these may be arranged in two differ- 
ent ways : the two resulting systems cannot be made to coincide by 
rotation, one being to the other as an object is to its mirror image. 




The carbon atom at the centre of the tetrahedron is then said to be 
asymmetric. 

The two optically active modifications of lactic acid would corre- 
spond to the two arrangements, or configurations. 



OH 



COOKN 




CH 3 



COOH N - N 



CH3 




OH 



Inactive lactic acid may be resolved into the active varieties, and 
must be regarded as resulting from the combination of equal propor- 
tions of these : the rotary power of the one would then be compensated 
by that of the other. 

When a molecule contains two asymmetric carbon atoms, the theory 
predicts four modifications of the compound. In the case of tartaric 
acid, COOH.CH(OH).CH(OH).COOH, the four forms are actually 
known. There are two active varieties, dextrotartaric acid and levo- 
tartaric acid, each of which contains two exactly similar asymmetric 
carbon atoms ; and of the two inactive modifications, mesotartaric acid 
contains two carbon atoms of opposite rotary power, while racemic 
acid (which can be split up into the dextro- and levo-forms) is com- 
posed of equal proportions of the two active tartaric acids. 



348 APPENDIX. 

This kind of isomerism, which is attributed to a different relative 
arrangement of the atoms in space, is called stereoisomerism, and that 
branch of theoretical chemistry which seeks to determine these spacial 
relations, stereochemistry. 

Stereoisomerism is not confined to those substances which contain 
asymmetric carbon atoms : many cases are also known among the un- 
saturated carbon compounds, and in certain classes of bodies contain- 
ing nitrogen. 



INDEX. 



Acetates, 210. 

Acetone, 210. 

Acetylene, 186. 

Acid, acetic, C 2 H*0 2 , 194, 207. 

acrylic, C 3 H 4 2 /216. 

antimonic, 135. 

arsenic, 130. 

arsenious, H 3 As0 3 , 130. 

benzoic, C 7 H<H) 2 , 231. 

boric, H 3 B0 3 , 138. 

butyric, OH 8 2 , 210. 

carbolic, 226. 

carbonic, H 2 C0 3 , 156. 

chloric, HC10 3 , 67. 

cbloroplatinic, PtCl 4 .2HCl, 333. 

chromic, H 2 CrO, 326. 

citric, C 6 H 8 7 , 219. 

cyanic, HOCX, 171. 

cyanuric, X 3 C 3 3 H 3 , 172. 

digallic, C^H^, 233. 

ethylsulphuric, C 2 H 5 .HS0 4 , 203. 

formic, CH 2 2 , 168, 193, 207. 

gallic, C7H60 5 , 232. 

hydracrylic, C 3 H 6 3 , 216. 

hydrazoic, X 3 H, 100. 

hydriodic, HI, 71. 

hydrobroinic, HBr, 69. 

hydrochloric,. HC1, 62. 

hydrocyanic, HCX T , 166. 
tests for, 167. 

hydrofluoric, HF, 71. 

hydrofluosilicic, H 2 SiF 6 , 142. 

hypochlorous, HC10, 65. 

hypophosphorous, H 3 P0 2 , 124. 

hyposulphurous, H 2 S 2 3 , 81. 

lactic, C 3 H 6 3 , 216. 

lauric, C 12 H 2 *0 2 , 215. 



Acid, malic, C 4 H 6 5 , 218. 
manganic, H 2 MnO*, 323. 
metaboric, HBO 2 , 139. 
metantimonic, HSbO 3 , 135. 
metaphosphoric, HPO 3 , 127. 
metarsenic, HAsO 3 , 130. 
myristic, C 14 H 28 2 , 215. 
nitric, HXO 3 , 112. 
nitro-hydrochloric, 114. 
nitrous, HXO 2 , 108. 
oleic, C 18 H 34 2 , 213. 
ortharsenic, H 3 As0 4 , 130. 
orthophosphoric, H 3 P0 4 , 125. 
oxalic, C 2 H 2 0*, 216. 
palmitic, C 16 H 32 2 , 212. 
permanganic, HMnO*, 324. 
phosphoric, 125. 
phosphorous, H 3 P0 3 , 125. 
picric, C 6 H 2 (X0 2 ) 3 OH, 227. 
propionic, C 3 H 6 2 , 210. 
pyroantimonic, H 4 Sb 2 7 , 135. 
pyroarsenic, H 4 As 2 7 , 130. 
pyrogallic, C 6 H 3 (OH) 3 , 232. 
pyrophosphoric, H 4 P 2 7 , 127. 
salicylic, C6H±.OH(C0 2 H), 232. 
silicic, 142. 
stannic, 330. 
stearic, C^H^O 2 , 213. 
succinic, OH60 4 , 218. 
sulphocarbonic, H 2 CS 3 , 163. 
sulphuric, H 2 S0 4 , 82. 

fuming, H 2 S 2 7 , 81. 

molecular structure of, 85. 
sulphurous, H 2 SO s , 80. 
tannic, 233. 
tartaric, C 4 H 6 6 , 218. 
tetraboric, H 2 B*0 7 , 139. 
349 



350 



INDEX. 



Acid, thiocyanic, HSCN, 174. 

thiosulphuric, H 2 S 2 3 , 81. 

uric, C5H 4 N*0 3 , 239, 340. 

valeric, &R™0 2 , 211. 
Acids, 51, 64. 

dibasic, 88, 216. 

fatty, 210, 212. 

of carbon, 207. 
Acrolein, C 3 H*0, 210. 
Affinity, 13. 
Agate, 140. 
Air, 92. 

carbon dioxide in, 95. 

water in, 94. 
Alabaster, 89. 
Albite, 305. 
Albumin, 336. 

Albuminoid substances, 335. 
Alcohol, amyl, C5H n .OH, 198. 

benzyl, C 6 H5.CH 2 OH, 230. 

butyl, CW.OH, 198. 

ethyl, C 2 H5.0H, 193. 
absolute, 194. 

methyl, CH 3 .OH, 192. 

propyl, C 3 IF.OH, 197. 
Alcoholic beverages, 195. 
Alcohols, 192. 

diatomic, 199. 

primary, 197. 

secondary, 197. 

tertiary, 198. 

triatomic, 200. 
Aldehyde, C 2 H±0, 194, 206. 

salicylic, CSHiOH.CHO, 231. 
Aldehydes, 206. 
Aldoses, 222. 
Ale, 196. 
Alizarin, 190. 
Alkali, 246. 
Alkaloids, 236. 
Alloys, 242. 
Alumina, A1 2 3 , 303. 
Aluminium, 301. 

bronze, 302. 

chloride, A1C1 3 , 302. 



Aluminium hydroxide, Al(OH) 3 , 303. 

oxide, A1 2 3 , 303. 

silicates, 305. 

sulphate, A1 2 (S0 4 ) 3 , 303. 

tests for, 306. 
Alums, 304. 
Amalgams, 242. 
Amethyst, 140. 
Amides, 173. 
Amines, 173, 229. 
Ammonia, NH 3 , 97. 

anatysis of, 99. 

combustion of, 99. 
Ammonium alum, 304. 

amalgam, 103. 

carbonates, 161. 

chloride, NH±C1, 101. 

chloroplatinate, 333. 

compounds, 100. 

cyanate, NH*O.CN, 172. 

dichromate, (NH 4 ) 2 Cr 2 7 , 327. 

molybdate, (NH 4 ) 2 Mo0 4 , 128. 

nitrate, NH*.N0 3 , 104. 

oxalate, (NH*) 2 C 2 0*, 117. 

phosphomolybdate, 128. 

picrate, NH 4 .OC 6 H 2 (N0 2 ) 3 , 228. 

sulphate, (NH) 2 S0 4 , 102. 

sulphide, (NH*)*S, 102. 

sulphocyanate, NH*NCS, 175. 

sulphydrate, NH*SH, 102. 

thiocyanate, 175. 
Amygdalin, 231. 
Amyl acetate, C5H".C 2 H 3 2 , 212. 

alcohols, C5HU.OH, 198. 
Amylene, &R™, 185. 
Analysis, 33. 

of carbon compounds, 190. 
Anatase, 331. 
Anglesite, 91. 
Aniline, C6H5.NH 2 , 228. 

colors, 230. 
Anthracene, C 14 H 10 , 190. 
Anthracite, 145. 
Antichlor, 81. 
Antimony, 134. 



INDEX. 



351 



Antimony chlorides, 135. 

oxides, 135. 

trisulphide, 135. 
Apatite, 127. 
Aragonite, 160. 
Argol, 218. 
Argon, 94. 

Aromatic compounds, 189. 
Arsenic, 128. 

chloride, AsCl 3 , 129. 

disulphide, As 2 S 2 , 130. 

pentasulphide, As 2 S 5 , 130. 

pentoxide, As 2 5 , 130. 

tests for, 131. 

trioxide, As*0 6 , 129. 

trisulphide, As 2 S 3 , 130. 
Arsenic oxide, As 2 5 , 130. 
Arsenious oxide, AsK) 6 , 129. 
Asbolite, 319. 
Atacamite, 288. 
Atmosphere, 92. 
Atomic heats, 250. 

theory, 40. 

weights, determination of, 41. 
Atomicity, theory of, 71, 78, 86, 110, 
136, 163, 178, 187, 197, 205, 245, 
257, 265, 272, 287, 291, 302. 
Atropine. C 17 H 23 X0 3 . 239. 
Auric chloride. AuCl 3 , 300. 
Avogadro's law, 39. 
Azurite, 289. 

Baking powders, 159. 
Barium, 270. 

carbonate, BaCO 3 , 160. 

chloride, BaCl 2 , 270. 

dioxide, BaO 2 , 271. 

hydroxide. Ba(OH) 2 . 51. 

hypophosphite, 124. 

monoxide, BaO, 270. 

nitrate, Ba(N0 3 ) 2 , 118. 

sulphate, BaSO 4 , 89. 

sulphide, BaS, 270. 

tests for, 271. 
Bases, 51. 



Beer, 196. 
Bell-metal, 287. 
Benzaldehyde, C6H5CHO, 230. 
Benzene, C 6 H 6 , 187. 

derivatives, 226. 
Benzine, 182. 
Benzyl alcohol, C 6 H*.CH 2 OH, 230. 

chloride, C 6 H5.CH 2 C1, 231. 
Beryl, 277. 
Bessemer process, 312. 

basic, 313. 
Bismuth, 295. 

chloride, BiCl 3 , 296. 

nitrate, Bi XO 3 ) 3 , 296. 

oxide, Bi 2 3 , 296. 

sulphide, Bi 2 S 3 , 296. 

tests for, 297. 
Bitter almond oil, 231. 
Bituminous coal, 145. 
Blende, 278, 282. 
Blue vitriol, CuSO*, 90. 
Bone-black, 148. 
Borax, Xa 2 BW, 139. 
Borneol, C 10 !! 1 ^, 235. 
Boron, 137. 

chloride, BC1 3 , 138. 

crystallized, 138. 

oxide, B 2 3 . 137. 

tests for, 140. 
Brandy, 196. 
Brass, 287. 
Braunite, 322. 
Britannia metal, 329. 
Bromides, 248. 
Bromine, 68. 
Bromoform, CHBr 3 , 206. 
Bronze, 287. 
Brookite, 331. 
Brucine. 240. 
Bunsen burner, 31. 
Butyl alcohols, C*H 9 .OH, 198. 
Butylenes, OH 8 , 185, 186. 

Cacodyl, As 2 (CH 3 )*, 210. 
Cadmium, 283. 



352 



INDEX. 



Cadmium ferrocyanide, Cd 2 (FeC 6 N 6 ), 
283. 

iodide, Cdl 2 , 283. 

oxide, CdO, 283. 

sulphide, CdS, 283. 

tests for, 283. 
Caesium, 257. 
Caffeine, C 8 Hi°N 4 2 , 238. 
Calcite, 160. 
Calcium, 265. 

acetate, 208. 

carbide, CaC 2 , 269. 

carbonate, CaCO 3 , 160. 

in hard water, 49, 156. 

chloride, CaCl 2 , 265. 

citrate, 220. 

hydroxide, Ca(OH) 2 , 267. 

hypochlorite, Ca(ClO) 2 , 66. 

lactate, Ca(C 3 H50 3 ) 2 , 216. 

phosphates, 126. 

phosphide, 123. 

sulphate, CaSO*, 89. 
in hard water, 49. 

tests for, 269. 
Calomel, Hg 2 Cl 2 , 292. 
Camphor, 234. 
Camphors, 234. 
Cane sugar, C 12 H 22 O n , 222. 
Caramel, 222. 

Carbamide, CO(NH 2 ) 2 , 173. 
Carbohydrates, 220. 
Carbon, 144. 

atomicity of, 163. 

compounds, analyses of, 190. 

dioxide, CO 2 , 153. 
in air, 95. 
tests for, 156. 

disulphide, CS 2 , 162. 

hydrates of, 220. 

monoxide, CO, 150. 

compounds of, 171. 

oxysulphide, COS, 164. 

reduction by, 150. 
Carbonates, 156. 

test for, 157. 



Carbonyl amide, CO(NH 2 ) 2 , 173. 

chloride, COC1 2 , 153. 

compounds of, 171. 
Carboxyl, 207. 
Case-hardening, 314. 
Casein, 338. 
Cassiterite, 328. 
Cast iron, 311. 
Caustic soda, 253. 
Celestite, 89, 269. 
Celluloid, 226. 
Cellulose, (C6H 10 O 5 ) n , 225. 

nitro-, 223. 
Cement, 267. 
Cerite, 306. 
Cerium, 306. 
Cerusite, 118, 272. 
Chalcedony, 140. 
Chaicocite, 284. 
Champagne, 196. 
Charcoal, 147. 

absorbent properties of, 149. 

animal, 147. 

filter, 149. 
Cheese, 339. 
Chemical affinity, 13. 

changes, 8. 

combination, 11. 

energy, 152. 

equations, 42. 

formulas, 42. 

laws and theories, 35, 38. 

nomenclature, 50, 61, 6Q. 

notation, 42. 
Chloral, C 2 C1 3 H0, 207. 
Chlorates, 66. 
Chlorides, 61, 248-. 

test for, 62. 
Chlorine, 57. 

analogies with bromine and 
iodine, 71. 
Chloroform, CHC1 3 , 206. 
Chromates, 326. 

test for, 328. 
Chrome green, 325. 



INDEX. 



353 



Chrome yellow, 327. 
Chromite, 325. 
Chromium, 325. 
alum, 304. 

chlorides, 325. 

oxides, 325. 

tests for, 327. 
Chromyl chloride, Cr0 2 Cl 2 , 328. 
Cinchona bark, 239. 
Cinchonine, C 19 H 22 X 2 0, 240. 
Cinnabar, 290, 294. 
Clay, 305. 
Cleveite, 245. 
Coal, 145. 

-mine explosions, 176. 

tar, 187. 
Cobalt, 319. 

blue, 320. 

chloride, CoCl 2 , 319. 

oxides, 319. 

tests for, 320. 
Cobaltite, 319. 
Cocaine, C"H«NO*, 239. 
Codeine, C^H^NO 8 , 239. 
Colcothar, 317. 
Collodion, 226. 
Combination, laws of, 35, 38. 
Combustion, 26. 

slow, 31. 
Compound ammonias, 236. 
Compounds, 9. 
Conine, C»H»N, 237. 
Copper, 2S3. 

acetate, Cu(C 2 H 3 2 ) 2 , 210. 

action of ammonia, 287. 

alloys of, 287. 

arsenite, CuHAsO 3 , 132. 

atomicity of, 287. 

carbonates, 289. 

chlorides, 287. 

matte, 285. 

metallurgy of, 284. 

electrolytic process, 285. 
wet process, 286. 

nitrate, Cu(N0 3 ) 2 , 118. 



Copper oxides, 289. 

pyrites, 284. 

sulphate, CuSO*, 90. 

sulphides, 289. 

tests for, 289. 
Copperas, FeSO*, 90. 
Corrosive sublimate, HgCl 2 , 292. 
Corundum, Al 2 3 , 303. 
Cream, 339. 
Cream of tartar, 219. 
Cresols, 230. 
Crocoite, 325. 
Cryolite, 72. 
Crystallization, 341. 

water of, 47. 
Crystallography, 341. 
Cupellation, 259. 
Cupric chloride, CuCl 2 , 288. 

ferrocyanide, Cu 2 (FeC 6 N 6 ), 289. 

nitrate, Cu(N0 3 ) 2 , 118. 

oxide, CuO, 288. 

sulphate, CuSO 4 , 90. 

sulphide, CuS, 288. 
Cuprite, 288. 
Cuprous chloride, Cu 2 Cl 2 , 288. 

oxide, Cu 2 0, 288. 

sulphide, Cu 2 S, 289. 
Cyamelide, 172. 
Cyanogen, (CN) 2 , 164. 

chloride, (CX) 3 C1 3 , 172. 

molecular structure of, 16b. 
Cymene, C^H 1 *. 234. 

Dalton's law, 36. 
Daturine, C^H^NO 8 , 239. 
Decomposition, 11. 

double, 15. 
Definite proportions, laws of, 36. 
Dextrin, 224. 
Dextrose. See Glucose. 
Dialysis, 142. 
Diamond, 144. 
Diastase, 222, 339. 
Didyrnium, 306. 
Digestion, 339. 



23 



354 



INDEX. 



Dimethylamine, (CH 3 ) 2 HN, 229. 
Dimorphism, 75, 343. 
Disaccharides, 223. 
Dolomite, 247. 
Dynamite, 200. 

Earthy cobalt, 319. 
Elements, 9. 

table of, 44. 
Elutriation, 282. 
Emerald, 277. 
Emery, 303. 
Epsom salt, MgSO 4 , 90. 
Equivalent combining proportions, 

38. 
Erbium, 306. 
Ethane, C 2 H6, 179. 
Ether, (C 2 H&) 2 0, 201. 
Ethers, compound, 211. 

simple, 201. 
Ethyl acetate, C 2 H5.C 2 H 3 2 , 211. 

bromide, C 2 H&Br, 205. 

formate, C 2 H5.0H0 2 , 212. 

hydrate, C 2 H5.0H, 193. 

iodide, C 2 H&L 204. 

nitrate, C 2 H&.N0 3 , 212. 

oxide, (C 2 H5) 2 0, 201. 

valerate, C 2 H5.C 5 H 9 2 , 212. 
Ethylene, C 2 H*, 184. 

bromide, C 2 H 4 Br 2 , 185. 

chloride, C 2 H 4 C1 2 , 184. 

cyanide, C 2 H 4 (CN) 2 , 218. 

hydrate, C 2 H 4 (OH) 2 , 199. 

oxide, C 2 H±0, 204. 
Ethylidene bromide, C 2 H*Br 2 , 206. 

Fats, natural, 213. 
Fehling's solution, 221. 
Feldspar, 305. 
Fermentation, acetic, 209. 

alcoholic, 193. 

lactic, 216. 
Ferric chloride, FeCl 3 , 316. 

ferrocyanide, Fe*(FeC 6 N ) 3 , 170. 

oxide, Fe 2 3 , 317. 



Ferric sulphate, Fe 2 (S0 4 ) 3 , 318. 

thiocyanate, 174. 
Ferrocyanides, 169. 
Ferromanganese, 311. 
Ferrous carbonate, FeCO 3 , 160, 307. 

chloride, FeCl 2 , 316. 

ferricyanide, Fe 3 (FeC 6 N 6 ) 2 , 171. 

oxide, FeO, 316. 

sulphate, FeSO±, 90. 
Fibrin, 337. 
Fire, 29. 
Fire-damp, 176. 
Fireworks, 271. 
Flame, 30, 176. 
Fluorine, 71. 
Fluor-spar, CaF 2 , 72. 
Formates, 208. 
Formulas, chemical, 42. 
Fractional distillation, 187. 
Fructose, C 6 H 12 6 , 221. 
Fuller's earth, 305. 
Fulminates, 195. 
Fusible metal, 296. 

Gadolinite, 306. 
Galactose, C 6 H 12 6 , 223. 
Galena, 272. 
Gallium, 300. 
Garnet, 305. 
Gas carbon, 147. 

illuminating, 145. 

liquor, 101, 147. 
Gases, manipulation of, 20. 

molecular volumes of, 39. 
Gasoline, 182. 
Gay-Lussac's laws, 36, 38. 
Gelatin, 339. ' 
Gelatinoid substances, 335. 
Germanium, 331. 
German silver, 287, 321. 
Gilding, 300. 
Gin, 197. 
Giobertite, 160. 
Glass, 141. 

etching on, 71. 



INDEX. 



355 



Glucinum, 277. 
Glucose, C6H1206, 193, 220. 
Glue, 339. 
Gluten, 223, 336. 
Glycerol, C 3 H*(OH) 3 , 199. 

ethers of, 213. 
Glycol, C 2 H±(OH)2, 199. 
Glycols, 198. 
Goethite, 307, 317. 
Gold, 297. 

assay, 300. 

chlorides, 299. 

oxides, 300. 
Graphite, 144. 
Greenockite, 283. 
Green vitriol, FeSO 4 , 90. 
Gum arabic, 225. 
Gums, 224. 

Gum tragacanth, 225. 
Gun-cotton, 225. 
Gun metal, 287. 
Gunpowder, 117. 
Gypsum, 89. 

Hausmannite, 322. 
Heat of combustion, 152. 
Heavy spar, BaSO*, 89, 270. 
Helium, 245. 
Hematite, 307. 
Hemoglobin, 337. 
Holmium, 306. 
Homologous bodies, 181. 
Horn silver, 261. 
Hydrates, 51, 246. 

of carbon, 220. 
Hydrazine, N 2 H 4 , 100. 
Hydrocarbons, C n H2n+2, 178. 

nomenclature of, 180, 183. 

unsaturated, 183. 

OH* 185. 
Hydrogen, 16. 

absorption by palladium, 22, 
334. 

antimonide, SbH 3 , 136. 

arsenide, AsH 3 , 132. 



Hydrogen, conductivity for heat, 21. 

diffusion of, 19. 

dioxide, 55. 

phosphide, PH 3 , 122. 

sulphide, H 2 S, 75. 
analysis of, 77. 
as reagent, 77. 
Hydroxide, 51, 246. 
Hydroxyl, 65. 
Hypochlorites, 66. 
Hypochlorous oxide, C1 2 0, 65. 

Iceland spar, 160. 
Illuminating gas, 145. 
Indigo, C16H10X2O 2 , 235. 

white, C^H^^O 2 , 236. 
Indium, 300. 
Ink, 233. 

sympathetic, 320. 
Iodides, 248. 
Iodine, 69. 

test for, 71. 
Iodoform, CHI 3 , 206. 
Iridium, 334. 
Iron, 307. 

blast-furnace process, 308. 

bloom, 310. 

carbonate, FeCO 3 , 160. 

cast, 311. 

Catalan process, 307. 

chlorides, 316. 

galvanized, 281. 

gray, 311. 

oxides, 316, 317. 

passive, 316. 

pig, 310. 

pyrites, 74, 307, 318. 

soft, 315. 

sulphates, 90, 318. 

sulphides, 317. 

tests for, 318. 

white, 311. 
Isomerism, 172, 186, 197, 345. 
Isometric system, 242. 
Isomorphism, 90, 343. 



356 



INDEX. 



Jet, 147. 

Kaolin, 305. 
Kerosene, 182. 
Ketoses, 222. 
Kieserite, 277. 
Kupfernickel, 320. 

Labradorite, 305. 

Lactose, C 12 H»O u , 223. 

Lamp-black, 148. 

Lanthanum, 306. 

Laughing-gas, N 2 0, 104. 

Law of Avogadro and Ampere, 39. 

of definite proportions, 36. 

of Gay-Lussac, 36, 38. 
Lead, 272. 

acetate, Pb(C 2 H 3 2 ) 2 , 210. 

carbonate, PbCO 3 , 160. 

chloride, PbCl 2 , 274. 

chromate, PbCrO*, 327. 

cupellation of, 259. 

dioxide, PbO 2 , 275. 

iodide, Pbl 2 , 274. 

monoxide, PbO, 275. 

nitrate, Pb(N0 3 ) 2 , 118. 

poisoning by, 274. 

red oxide, Pb 3 4 , 275. 

sulphate, PbSO*, 91. 

sulphide, PbS, 276. 

tests for, 276. 
Legumine, 336. 
Lepidolite, 250. 
Levulose, C 6 H 12 6 , 221. 
Life, chemistry of, 334. 
Lignite, 147. 
Lime, CaO, 266. 

chlorinated, CaCl(ClO), 66, 268. 
Limestone, 160. 
Limonite, 307. 
Litharge, 275. 
Lithium, 250. 
Lixiviation, 159. 
Lunar caustic, 118. 

Magenta, 230. 



Magnesia, MgO, 278. 

alba, 160. 
Magnesite, 160, 277. 
Magnesium, 277. 

ammonium phosphate, 278. 
carbonate, MgCO 3 , 160. 
chloride, MgCl 2 , 277. 
citrate, 220. 

hydroxide, Mg(OH) 2 , 278. 
oxide, MgO, 278. 
sulphate, MgSO*, 90. 
tests for, 278. 
Magnetite, 307. 
Malachite, 289. 
Malt, 193. 

Maltose, C 6 H 12 6 , 196, 224. 
Manganese, 322. 

dioxide, MnO 2 , 322. 
oxides, 322. 
sulphide, MnS, 324. 
tests for, 324. 
Marble, 160. 
Marcasite, 318. 
Marl, 305. 
Marsh gas, 178. 
Marsh's test for arsenic, 133. 
Massicot, 275. 
Matches, 121. 
Meadow-sweet oil, 231. 
Menthol, C 10 H 2 °0, 235. 
Mercuric chloride, HgOl 2 , 292. 
cyanide, Hg(CN) 2 , 169. 
iodide, Hgl 2 , 293. 
nitrate, Hg(N0 3 ) 2 , 118. 
oxide, HgO, 293. 
Mercurous chloride, Hg 2 Cl 2 , 292. 
iodide, Hg 2 I 2 , 293. 
oxide, Hg 2 0, 293. 
Mercury, 290. 

atomicity of, 291. 
chloride, 292. 
cyanide, 169. 
fulminate, 195. 
molecular weight of, 291 
oxides, 293. 






INDEX. 



357 



Mercury, sulphide, 294. 

tests for, 294. 
Metallic bromides, 248. 

carbonates, 156. 

carbonyls, 153. 

chlorides, 61, 248. 

hydrates, 52, 246. 

nitrates, 115, 116. 

oxides, 246. 

sulphates, 87. 
Metals, 15. 

general properties of, 241. 

natural state of, 242. 
Methane, CH 4 , 175. 
Methylamine, CH 3 .NH 2 , 229. 
Methylaniline, CH 3 .C6H6N, 229. 
Methylbenzene, CH 3 .C 6 H&, 139. 
Methyl chloride, CH 3 C1, 204. 

cyanide, CH 3 CN, 210. 

hydroxide, CH 3 .OH, 192. 

iodide, CH 3 I, 178. 

oxide, (CH 3 )*0, 201. 

salicylate, CHWIIH) 8 , 232. 
Mica, 305. 
Milk, 338. 
Mineral waters, 49. 
Minium, 275. 
Mispickel, 128. 
Molecular weights, determination of, 

40. 
Molecules, 11. 
Molybdenite, 328. 
Molybdenum, 328. 
Monazite, 306, 331. 
Monoclinic system, 343. 
Monosaccharides, 221. 
Morphine, C"H 19 N0 3 , 329. 
Myosin, 337. 

Naphtha, 182. 
Naphthalene, C 10 H8, 189. 
Xarcotine, C 22 H 23 N0 7 , 239. 
Native metals, 242. 
Neodymium, 306. 
Nessler's reagent, 293. 



Niccolite, 320. 
Nickel, 320. 

carbonyl, Ni(CO) 4 , 153. 

chloride, NiCl 2 , 322. 

oxides, 322. 

plating, 321. 

sulphate, NiSO 4 , 322. 

tests for, 322. 
Nicotine, C^H^N 2 , 238. 
Niobium, 136. 
Nitrates, 115, 116. 
Nitric oxide, NO, 106. 
Nitrobenzene, C 6 H 5 .N0 2 , 228. 
Nitrogen, 91. 

atomicity of, 110. 

bromide, NBr 3 , 104. 

chloride, NCI 3 , 104. 

dioxide, NO, 106. 

group of elements, 136. 

iodide, 103. 

monoxide, N 2 0, 104. 

pentoxide, N 2 5 , 110. 

peroxide, NO 2 , 108. 

trioxide, N 2 3 , 109. 
Nitroglycerin, C 3 H 5 (N0 3 ) 3 , 200. 
Nitrosyl chloride, NOC1, 108. 
Nitrotoluenes, C 6 H*(CH 3 )N0 2 , 229. 
Nitrous oxide, N 2 0, 104. 
Nomenclature of acids and salts, 66. 

of chlorine compounds, 61. 

of oxygen compounds, 50. 
Notation, 42. 

Oils, essential, 189. 

fatty and drying, 214. 
Olein, C 3 H5(C 18 H 33 2 ) 3 , 214. 
Opium, 239. 
Organic compounds, 173. 

chemistry, 173. 
Orpiment, As 2 S 3 , 130. 
Orthoclase, 305. 
Orthophosphates, 126. 
Orthorhombic system, 243. 
Osmium, 334. 
Oxalates, 217. 



358 



INDEX. 



Oxides, 50, 246. 
Oxygen, 23. 

in air, 92. 

manufacture of, 323. 

properties of, 26. 
Oxyhydrogen blow-pipe, 29. 
Ozone, 53. 

Palladium, 333. 

Palmitine, C^C^H^O 2 ) 3 , 213. 

Paraffin, 182. 

Paris green, 133. 

Pepsin, 339. 

Petroleum, 181. 

Pewter, 274. 

Phenol, (OT.OH, 226. 

nitro-, 227. 

test for, 227. 
Phosphorus, 119. 

amorphous, 121. 

chlorides, 123. 

oxides, 123. 
Phosphonium salts, 123. 
Photography, 264. 
Physical changes, 7. 
Pig iron, 310. 
Pitchblende, 328. 
Plaster of Paris, 89. 
Platinum, 332. 

black, 333. 

chlorides, 333. 

sponge, 332. 
Plumbago, 144. 
Polymerism, 185. 
Polysaccharides, 226. 
Porter, 196. 
Potassium, 254. 

acid carbonate, KHCO 3 , 160. 

acid tartrate, KC 4 H 5 6 , 219. 

alum, 304. 

antimonyl tartrate, 219. 

bromide, KBr, 256. 

carbonate, K 2 C0 3 , 159. 

chlorate, KC10 3 , 69. 

chloride, KC1, 256. 



Potassium chloroplatinate, 257, 333. 

chromate, K 2 O0*, 326. 

cyanate, KOCN, 171. 

cyanide, KCN, 168. 

dichromate, K 2 Cr 2 7 , 326. 

ferri cyanide, 170. 

ferrocyanide, K 4 FeC 6 N 6 , 169. 

hydroxide, KOH, 255. 

hypochlorite, KC10, 66. 

iodide, KI, 256. 

manganate, K 2 Mn0 4 , 323. 

nitrate, KNO 3 , 116. 

oxide, K 2 0, 254. 

permanganate, KMnO 4 , 324. 

picrate, KO.C 6 H 2 (N0 2 ) 3 , 228. 

-sodium tartrate, 219. 

sulphate, K 2 S0 4 , 89. 

sulphocyanate, KNCS, 174. 

sulphydrate, KSH, 78. 

tartrate, K 2 OH±0 6 , 219. 

tests for, 256. 

thiocyanate, 174. 
Pottery, 305. 
Praseodymium, 306. 
Propane, C 3 H 8 , 179. 
Propyl alcohols, C 3 H7.0H, 197. 
Propylene, C 3 H 6 , 185, 186. 
Prussian blue, 170. 
Ptyalin, 339. 
Purple of Cassius, 300. 
Pyrites, copper, 284. 

iron, 318. 
Pyrogallol, C«H 3 (OH) 3 , 232. 
Pyrolusite, 322. 
Pyroxylin, 225. 

Quartz, 140. 

Quinine, C 2 °H 2 *N 2 2 , 239. 
sulphate, 240. 

Radicals, 87. 

acid and basic, 103. 
hydrocarbon, 178. 

hydrates of, 192. 

oxides of, 201. 



INDEX. 



350 



Realgar, As 2 S 2 , 130. 

Red precipitate, 293. 

Reinseh's test for arsenic, 131. 

Respiration, 32. 

Rhodium, 334. 

Rhuthenium, 334. 

Rochelle salt, KNaC*H 4 06, 219. 

Rock crystal, 140. 

Rosaniline, C 2 <>H 2 iN 3 0, 229. 

Rubidium, 257. 

Ruby, 304. 

Rum, 197. 

Rust, 315, 317. 

Ruthenium, 334. 

Rutile, 330. 

Saccharose, C^fl^O 11 , 224. 
Safety-lamp, 176. 
Saltpetre, 116. 
Salts, 65. 

ethereal, 211. 

neutral and acid, 88. 
Samarium, 306. 
Sand, 140. 
Saponification, 214. 
Sapphire, 304. 
Scandium, 306. 
Scheele's green, 133. 
Sea-water, 253. 
Selenite, 89. 
Selenium, 87. 
Serpentine, 277. 
Shot, 129, 274. 
Siderite, 307. 

Siemens regenerative furnace, 279. 
Silica, SiO 2 , 140. 
Silicon, 140. 

oxide, SiO 2 , 140. 
Silver, 257. 

arsenate, Ag 3 AsO±, 130, 133. 

arsenite, Ag 2 HAs0 8 , 132. 

assay, 262. 

chloride, AgCl, 261. 

chromate, Ag 2 CrO*, 262. 

cyanide, AgCN, 168. 



Silver iodide, Agl, 262. 

nitrate, AgNO 3 , 118. 

orthophosphate, Ag3P0 4 , 126. 

oxide, Ag 2 0, 261. 

sulphide, Ag 2 S, 262. 

tests for, 262. 
Silvering, 262. 
Slow combustion, 31. 
Smalt, 320. 
Smaltite, 319. 
Smithsonite, 160, 278. 
Soap, 214. 

salt-water, 215. 
Soapstone, 277. 
Soda-water, 154. 
Sodium, 251. 

acetate, XaC 2 H 3 2 , 210. 

acid carbonate, NaHCO 3 , 158. 

acid sulphate, NaHSO 4 , 88. 

alum, 304. 

borates, 137, 139. 

carbonate, Na 2 C0 3 , 157. 

chloride, NaCl, 253. 

dioxide, Na 2 2 , 253. 

hydroxide, NaOH, 252. 

hypochlorite, NaCIO, 66. 

hyposulphite, Xa 2 S 2 3 , 81. 

methylate, NaCH 3 0, 193. 

nitrate, NaXO 3 , 116. 

oxide, Na 2 0, 253. 

phosphates, 126. 

potassium tartrate, 219. 

sulphate, Na 2 S0 4 , 88. 

sulphite, Na 2 S0 3 , 80. 

tests for, 254. 

tetraborate, Na 2 B 4 7 , 139. 

thiosulphate, Na 2 S 2 3 , 81. 
Solder, 274. 
Soluble glass, 142. 
Spathic iron, 160, 307. 
Specific heat, 249. 
Spectroscope, 244. 
Spectrum analysis, 242. 
Speculum metal, 287. 
Speiss, 320. 



360 



INDEX. 



Spiegel eisen, 311. 

Spinel, 247. 

Stannic chloride, SnCl 4 , 330. 

oxide, SnO 2 , 330. 
Stannous chloride, SnCl 2 , 330. 

oxide, SnO, 330. 
Starch, (C 6 Hi°05)a ? 223. 
Stearin, C 3 fl 5 (C 18 H**) 2 ) 3 , 213. 

candles, 214. 
Steel, 311. 

Bessemer process, 312. 

tempering, 313. 
Stereoisomerism, 345. 
Stereochemistry, 344. 
Stibnite, 134. 
Strontianite, 160, 269. 
Strontium, 269. 

carbonate, SrCO 3 , 160. 

chloride, SrCl 2 , 269. 

dioxide, SrO 2 , 270. 

hydroxide, Sr(OH) 2 , 270. 

monoxide, SrO, 270. 

nitrate, Sr(N0 3 ) 2 , 118. 

sulphate, SrSO*, 89. 

sulphide, SrS, 269. 

tests for, 270. 
Strychnine, C 2 1H 22 N 2 2 , 240. 
Substance, definition, 7. 
Sugar, cane, C 12 H 22 1J , 224. 

grape, C6Hi 2 6 , 223. 

milk, C 12 H 22 O n , 225. 

of lead, Pb(C 2 H 3 2 ) 2 , 210. 
Sulphates, 87. 

test for, 88. 
Sulphides, 73, 248. 

tests for, 75. 
Sulphites, 80. 

Sulpho-urea, CS(NH 2 ) 2 , 175. 
Sulphur, 73. 

atomicity of, 86. 

dimorphism of, 75. 

dioxide, SO 2 , 79. 

soft, 74. 

trioxide, SO 3 , 81. 
Sulphuryl chloride, S0 2 C1 2 , 86. 



Sulphydrates, 78. 
Symbols, 42. 
Synthesis, 33. 

Tannin, 233. 

Tanning, 233, 

Tantalum, 136. 

Tartar-emetic, K(SbO)C*HK) 6 , 219. 

Tartrates, 219. 

Tellurium, 87. 

Thallium, 300. 

Theine, C 8 Hi°N*0 2 , 238. 

Theobromine, C 7 H 8 N*0 2 , 238. 

Thiosulphates, 81. 

Thorite, 331. 

Thorium, 331. 

Thulium, 306. 

Thymol, C 10 H 14 O, 234. 

Tin, 328. 

dichloride, SnCl 2 , 329. 

oxides, 330. 

sulphides, 330. 

tests for, 331. 

tetrachloride, SnCl*, 330. 
Titanium, 331. 
Toluene, C6H5.CH 3 , 189. 
Topaz, 304. 

Trichloraldehyde, C 2 C1 3 H0, 207. 
Triclinic system, 344. 
Trimethylamine, (CH 3 ) 3 N, 229. 
Trinitrophenol, C6H 2 (N0 2 ) 3 OH, 227. 
Triphylite, 250. 
Tungsten, 328. 
Turnbull's blue, 171. 
Turpentine, C 10 H 16 , 189. 
Type-metal, 135. 

Ultramarine, 306. 
Uranium, 325. 
Urea, CO(NH 2 ) 2 , 172. 

molecular structure of, 173. 

nitrate, CO(NH 2 ) 2 HN0 3 , 174. 

Vanadium, 136. 
Vegetable parchment, 225. 



INDKX 



361 



Verdigris, 210, 286. 
Vermilion, 294. 
Vinegar, 209. 
Vitriol, blue, CuSO*, 90. 

green, FeSO, 90. 

oil of, IPSO*, 83. 

white, ZnSO*, 90. 

Water, 32. 

electrolysis of, 33. 

hard, 48. 

in air, 94. 

mineral, 49. 

natural, 48. 

of crystallization, 47. 

properties of, 45, 47. 

synthesis of, 34. 
Water-gas, 153. 
Whiskey, 196. 
White indigo, 236. 
White lead, 160. 
White vitriol, ZnSO 4 , 90. 



Wine, 196. 

Wintergreen oil, 232. 
Witherite, 160, 270. 
Wolfram, 328. 
Wood-spirit, 192. 
Wrought-iron, 310. 

Yeast, 194. 
Ytterbium, 306. 
Yttrium, 306. 

Zinc, 278. 

carbonate. ZnCO 3 , 160. 

chloride, ZnCl 2 , 281. 

ferrocyanide, Zn 2 (FeC 6 X 6 ), 282. 

oxide, ZnO, 282. 

sulphate, ZnSO, 90. 

sulphide, ZnS, 282. 

tests for. 282. 

white, ZnO, 2S2. 
Zircon, 331. 
Zirconium, 331. 



THE END. 



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